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COKRIGH.T DEPOSm 



A PRACTICAL CHEMISTRY 

FOR 
HIGH SCHOOL STUDENTS 



A PRACTICAL CHEMISTRY 

FOR 
HIGH SCHOOL STUDENTS 



BY 

CHARLES GILPIN COOK, Ph.D, 

OF THE BOYS* HIGH SCHOOL, NEW YORK CITY 




ILLUSTRATED 



D. APPLETON AND COMPANY 
NEW YORK CHICAGO 



OH 33 
.C7Z 



Copyright, 1914, 1916, by 
D. APPLETON AND COMPANY 



SEP 16 1916 
©CIA437716 



PREFACE 

This book is intended for those schools desiring a more 
practical and interesting course in chemistry. Its purpose 
is to bring the pupil into closer contact with the things of 
everyday life, to give him an insight into the nature of 
those chemical processes with which all of us come more or 
less in contact, to furnish him with some chemical knowl- 
edge of the things necessary for his health and comfort 
and thus better adjust him to his environment. 

In doing this no attempt is made to treat the subject in 
a theoretical way ; in fact, but little of the theory of chem- 
istry is mentioned. The idea of atoms and the use of 
symbols and formulae have, however, been introduced after 
a few substances have been studied. 

The mathematical side of chemistry finds no place in a 
text of this sort, but a chapter on chemical arithmetic has 
been included in the Appendix. 

The book is the result of classroom experiment and has 
grown out of an attempt in the Boys ' High School to adapt 
chemistry teaching to the needs of second-year pupils. 
The subjects chosen, however, have not been those which 
appeal particularly to boys, but are those which are of 
most vital importance to both girls and boys and, in gen- 
eral, are of equal interest to both. 

Charles G. Cook. 
New York City. 



CONTENTS 

CHAPTER PAGE 

I. Introduction . 1 

II. Water 5 

III. Oxygen . . ... . . .17 

IV. Hydrogen 25 

V. Acids, Bases and Salts 34 

VI. Atoms, Symbols, Formula and Equations . 41 

VII. Salt 49 

VIII. The Air ........ 59 

IX. Carbon and Fuels .69 

X. Flames . . . . . . . .87 

XL Phosphorus and Matches .... 96 

XII. Sulphur and Some of Its Compounds . . 102 

XIII. Compounds op Nitrogen and Explosives . 113 

XIV. Fertilizers 128 

XV. Food . .140 

XVI. Some Familiar Compounds Used in Medicine 175 

XVII. Fibers 191 

XVIII. Cleaning Materials 202 

XIX. Building Materials 212 

XX. Alloys 252 

XXI. Metals Used for Plating and Decorating . 259 

XXII. Glass and Pottery . . . . . . 268 

XXIII. Paints 279 

XXIV. Photography . 287 

Appendix A. — Experiments for Demonstration and 

Laboratory Exercises .... 295 

Appendix B. — Chemical Arithmetic .... 330 

Index 353 

vii 



LIST OF ILLUSTRATIONS 

FIG. PAGE 

1. — The chemist at work 2 

2. — Snow crystals photographed by W. A. Bentley . 5 

3. — Thermometers 6 

4. — Sectional diagram of a still for distilling water in 

small quantities 9 

5. — Explanation of sand filter 10 

6. — Sand filter at Washington, D. C, before the sand was 

put in 11 

7. — The electrolysis of water ...... 13 

8. — Generating oxygen 19 

9. — Generating a tank of oxygen . . . .21 

10. — Making ozone ........ 22 

11. — Hydrogen not a supporter of combustion ... 25 

12. — Pouring hydrogen ....... 26 

13. — Burning air in illuminating gas 27 

14. — The oxy-hydrogen blowpipe 28 

15. — Water from burning hydrogen 29 

16. — Decomposition of steam by heated iron ... 29 

17. — Generating and collecting hydrogen .... 31 

18. — Neutralization 38 

19. — Electrolysis of sodium chloride . . . .50 

20. — Apparatus for the manufacture of sodium hydroxide 

by electrolysis 51 

21. — Apparatus for preparation of sodium by electrolysis : 

Castner's process 52 

22. — Generating and collecting chlorine .... 54 

23. — Diagram of bleaching . 55 

24. — The principle of the barometer 60 

25. — The volume of oxygen in air 61 

26. — Combustion produces carbon dioxide.. ... 62 

ix 



x LIST OF ILLUSTRATIONS 



PAGE 



27. — Action of green plants upon carbon dioxide . . 65 

28. — Graphite crucibles . 70 

29. — Collecting marsh gas from the bottom of a pond . 75 

30. — Apparatus for generating acetylene .... 76 

31. — The acetylene burner ....... 77 

32. — The acetylene torch in use 78 

33. — Diagram of the manufacture of water gas . . 82 

34.— The blast lamp . 88 

35. — Reducing and oxidizing flames 89 

36. — A gas mantle . 89 

37. — Manufacturing gas mantles 90 

38. — Drying the mantles . 91 

39. — A gauze placed over a burner with the escaping gas 

lighted above the gauze 92 

40. — Davy's safety-lamp 92 

41. — Fire extinguisher 93 

42. — Phosphorous furnace 97 

43. — Sulphur flowing from a sulphur well .... 103 

44. — Many thousands of tons of sulphur in one pile . . 104 

45. — Diagram of a sulphuric acid plant .... 106 

46. — The Glover tower in a sulphuric acid plant . . 107 

47. — The Gay-Lussac tower in a sulphuric acid plant . 108 

48. — Making nitric acid . . . . . . . 114 

49.— Making ice . .118 

50.— A blast 122 

51.— An unfertilized field . . ... . . .128 

52. — A field fertilized with potash 129 

53. — More complete fertilization of the same field . . 130 

54.— A fertile field 131 

55. — Nodules on the roots of clover 133 

56. — A demonstration of making nitric acid by the electric 

arc 134 

57. — Diagram showing transformations of nitrogen in 

nature 135 

58. — Charts showing composition of various foods . . 141 

59. — A diagram showing section of a vacuum pan . . 143 

60. — A modern cow stable 150 

61. — The pasteurizing room of a large city milk dealer . 152 



LIST OF ILLUSTRATIONS xi 

FIG. PAGE 

62. — Many thousands of clean bottles ready to receive the 

pasteurized milk 153 

63. — The De Laval milk separator 155 

64. — Butter-making. The old way 157 

65. — Butter-making. The new way 157 

66. — Making cider 159 

67. — Tanks in which the cider changes to vinegar . . 159 

68. — Leaves, flowers and pods of the cocoa plant . . 165 

69. — Cocoa pods containing the cocoa beans . . . 166 

70. — Moulding cocoa 167 

71.— A cotton field in Texas 192 

72.— A cotton gin . . 193 

73.— The fiber industry 195 

74.— Silk cocoons 196 

75. — A spinning machine 197 

76. — Cutting large blocks of soap 203 

77. — Various processes in the manufacture of shaving 

soap 204 

78.— Building a high school 213 

79. — A marble quarry in Vermont 214 

80. — Diagram showing the shape of perfect gypsum crystals 217 
81. — The proportion in which the cement, sand and stone 

should be mixed in making concrete . . . 219 

82. — Mixing concrete by hand 219 

83. — A wood frame in which to build a concrete wall . 220 

84.— The Woolworth Building 222 

85. — Diagram of cross-section of a blast furnace . . 223 

86. — The Bessemer converter 226 

87. — The converter in action 228 

88. — The open-hearth furnace 229 

89. — Hammer for making drop-forgings .... 230 

90.— A lead mine 232 

91. — A reverberatory furnace 233 

92. — At the bottom of a lead blast furnace . . . 234 

93. — A machine for casting pig lead 234 

94. — Making lead pipe 235 

95. — In a lead pipe factory 236 

96.— Rolling sheet lead 237 



xii LIST OF ILLUSTRATIONS 

FIG. PAGE 

97. — Diagram showing the refining of copper by the elec- 
tric current . . . . . . . 241 

98. — Hall's redaction furnace for aluminium . . . 243 
99. — Use of thermit in welding the stern shoe of a steam- 
ship . . 244 

100. — Carborundum furnace 245 

101. — Cross-section of carborundum furnace . . . 246 

102. — Carborundum process 247 

103A.-B.— Fusible links 254 

104. — Linotype machine, in which type metal is cast into 

type by pressing the letters on the keyboard . . 255 
105. — Stamping money in the United States Mint, Philadel- 
phia . . . . . . . . 256 

106.— Hydraulic mining 260 

107. — Desilvering lead . . . . ' . . . . 264 

108.— Silver plating 265 

109. — Exhibit of the process of making glass . . . 269 

110. — Machine for blowing window glass .... 272 

111. — Filter-press for filtering clay after washing . . 273 
112. — Pressing tiles from the clay which has been removed 

from the filter-press shown in Figure 111 . . 274 

113. — Kilns for burning tiles 275 

114. — Tile floor in barber shop of Woolworth Building, 

New York City 276 

115. — White lead pot and buckle 280 

116.— Making white lead . . . . . . .280 

117.— Filling the white lead pots 281 

118. — Expressing linseed oil . . . . . . . 282 

119A.— A negative 290 

119B. — Print, or positive corresponding to the negative in 

Figure 119A 290 

120.— The hydrogen gun . . . . . . .300 

121. — Reduction by charcoal ...... 307 

122. — Collecting a test tube full of ammonia . . 312 

123. — Fermentation of glucose 314 

124. — Distillation of alcohol from the fermented mash . 315 

125. — Redistillation of alcohol 316 

126.— Babcock milk testing bottle 317 



LIST OF ILLUSTRATIONS xiii 

FIG. PAGE 

127.— Babcock milk tester 318 

128.— Apparatus to illustrate the Solvay process . . 319 

129.— Making ether 322 

130. — Liberating bromine 323 

131.— Test for carbonates .327 



A PRACTICAL CHEMISTRY FOR 
HIGH SCHOOL STUDENTS 



CHAPTER I 
INTRODUCTION 

Children are fond of pulling things apart ; perhaps they 
wish to try to see how things are made and of what 
they are made. Older people are in some respects like 
children; they also want to know how things are made 
and of what they are made. But the objects of their 
curiosity are not, as a rule, articles manufactured by 
other men, but rather the vast number of things which 
are found in nature, such as water, air, wood, stone, 
coal, flour, fruits, eggs, etc. Grown people have done 
much at this taking things apart and have found out 
much about the composition of many things in nature 
and how they are made. They have called the process of 
taking things apart to learn their composition analysis. 
In many cases also men have learned to put things to- 
gether to make substances and this is called synthesis. 
The people who do this kind of work are said to be 
chemists, and the study connected with the work is called 
chemistry. It is a very large study and its importance 
is very great, for it has done much to improve many of 
our manufactures. 

Physical and Chemical Properties and Physical and 
Chemical Changes. — When a new toy is given to a little 

1 



2 A PRACTICAL CHEMISTRY 

boy he does not generally break it up right away to see 
where "the squeak comes from," but he looks at it and 
tries to see what he can do with it in every way he can 
think of. So the chemist, also, first of all carefully ex- 
amines any new substance to learn its physical properties, 
that is, what qualities it will show without being treated 




Fig. 1. — The Chemist at Work. The interior of a laboratory at a leadworks. 

with acids and other chemicals. Later he studies the 
behavior of the new substance when in contact with other 
substances, and the properties thus exhibited are chemical 
properties. Now the chemist in bringing out the physical 
properties of a substance does not change its composition, 
but in showing its chemical properties he does change 
the composition. All changes, then, which do not change 
the composition of a substance are said to be physical 



INTRODUCTION 3 

changes, while those which do change the composition of 
a substance are called chemical changes. You may think 
of many examples of these changes : when a tree is blown 
down by a storm, when it is cut into firewood, or when 
it is made into lumber, all of these are physical changes, 
for the composition of the wood is not changed. On the 
other hand, if the wood is burned in the fire or rots 
through exposure to the weather, its composition is 
changed and the change is chemical. 

Elements. — When the small boy has torn his toy apart 
he has left bits of wire, cloth, wood, sawdust, etc., and 
these he is unable to break up into anything more simple, 
and, if he has found out all he wished to know, he is 
satisfied. Likewise the chemist finds that he can go only 
so far in his analysis, and, like the small boy, comes to 
substances which he can tear apart no further. These 
substances he calls elements. You have often heard the 
names of some of them, as iron, lead, sulphur, gold, silver, 
carbon. There are about eighty of these elements known, 
but many of them are rare. 

What Chemistry Teaches. — You will understand, then, 
that, in our study of chemistry, we wish to learn these 
elements, in what substances they are found, how we can 
get them, what they will do, what laws govern them. 

SUMMARY 

Curiosity has led to many discoveries. The composition of sub- 
stances is ascertained by both analysis and synthesis. 

Analysis — Analysis is the separating of a substance into its con- 
stituents for the purpose of finding its composition. 

Synthesis. — Synthesis is the building up of a substance from its 
constituents. 

Physical Properties. — The physical properties of a substance are 
those which may be discovered without changing the composi- 
tion of the substance. 



4 A PRACTICAL CHEMISTRY 

Physical Changes. — Physical changes are those which do not 

change the composition of the substance. 
Chemical Properties. — The chemical properties of a substance are 

those made manifest by its behavior when in contact with 

other substances. 
Chemical Changes.— Chemical changes are those which change the 

composition of the substance. 
Elements. — Elements are substances which cannot be decomposed 

into any simpler form of matter by known means. 
Chemistry. — The science of chemistry teaches the sources, proper- 
ties, and uses of the elements and the laws which govern 

them. 

REVIEW QUESTIONS 

1. Why is curiosity sometimes desirable? 

2. What is analysis? 

3. What is synthesis? 

4. Germany has many chemists; what effect has this had on 
the development of that country? 

5. What is a physical change? 

6. A chemical change? 

7. Is the rusting of iron a chemical or a physical change? 

8. What are elements? 

9. Why is it possible that some of the substances now known 
as elements may not be elements? 

10. What does chemistry teach? 



CHAPTER II 



WATER 



General Properties. — We shall begin our study with 
water. Perhaps there are some facts about this very- 
familiar substance which will be new to you. "What be- 
comes of water when it dries, or evaporates, as we some- 
times say? If our eyes were sharp enough we might see 
it flying off into the air as little particles. If the air 
rises to a colder part of the sky, the little particles run 




Fig. 2. — Snow Crystals Photographed by W. A. Bentley. Mr. Bentley has made 
many hundreds of pictures of snow crystals and among these no two are alike. 
Perhaps no two snow crystals ever were alike. They are all, however, symmetrical 
in form and each has six rays. The student will find the study of snow a fascinating 
pastime. 



6 



A PRACTICAL CHEMISTRY 



100 



2J2 



together to form drops and may fall as rain. Sometimes 
the water vapor in the air (called moisture) freezes be- 
fore it can form raindrops, and then we have snowflakes. 
If snowflakes are examined, it will be seen that they all 
have six points. (See Fig. 2.) 

Pitchers often crack when water freezes in them. The 
water increases in volume when it freezes ; if you should 
c F measure carefully, you would find it 

expands about one-tenth of its volume. 
You might also find that hot water 
while cooling gradually grows smaller, 
or contracts, until the ■ temperature of 
four degrees centigrade (written 4° 
C.) is reached. It then begins to ex- 
pand and continues to expand till 
frozen. Water is therefore heaviest at 
4° C. At this temperature its density 
is said to be 1. 

Now just a word about the ther- 
mometer with which we measure tem- 
perature. Fahrenheit thermometers 
are used at home, but in chemistry 
centigrade thermometers are used. 
The two kinds differ only in the figures 
and marks upon them. (See Fig. 3.) 
If both thermometers are put into 
melting ice, the mercury will come to 
the 0° mark on the centigrade, and to 
the 32° mark on the Fahrenheit. Like- 
wise, if they are put into steam over freely boiling water, 
the temperature of the steam will be found to be 100° 
centigrade and 212° Fahrenheit. On the centigrade ther- 
mometer there are 100 degrees between the freezing and 
boiling points of water, while on the Fahrenheit there are 
180 degrees. 



32 



§ § 



Fig. 3. — Thekmometers. 
It will be noticed that 
0° Centigrade corre- 
sponds to 32° Fahren- 
heit and likewise 100° 
C. to 212° F. 



WATER 7 

Water is used as a standard for temperatures. It is 
also a standard for a number of other measurements. 
We compare the weight of other objects with water. If, 
for instance, we find a substance seven times as heavy as 
water, we say it has a specific gravity of 7. In the same 
way we compare the weights of liquids with water ; if a 
liquid is said to have a specific gravity of 1.8, we mean 
it is 1.8 times as heavy as water. 

Solution. — Perhaps you have noticed that when sugar 
or salt is dropped into water it slowly disappears. What 
becomes of it? It has not been destroyed, for when the 
water dries away the salt or sugar is left behind. What 
happens is that the substance simply separates into little 
particles, called molecules, and these move about all 
through the water and are said to form a solution with 
the water. Many other substances act in the same way, 
that is, dissolve when put into water. Nearly all solids 
dissolve faster and in larger quantities in hot water than 
in cold, while gases dissolve in larger quantity in cold 
water and under pressure. You have noticed how the 
gas escapes from soda water when drawn and allowed 
to stand in a glass. This was because it was under pres- 
sure and this pressure was removed when the soda came 
into the glass. If you let cold water stand in a glass in 
a warm room, you will soon find little bubbles of air 
sticking to the inside of the glass. This air was dissolved 
in the cold water, but could not remain in solution after 
the water became warm. The molecules of gas move so 
fast that they cannot remain in the warm water; if the 
water is cooled or if the gas is put under pressure, the 
molecules will not move so fast, and therefore more gas 
will go into solution under these conditions. When a solid 
dissolves its molecules must move faster, and we may 
cause them to do this by heating the solution. If a sub- 
stance dissolves without heating, the solution is often 



8 A PRACTICAL CHEMISTRY 

found to grow cool, since the molecules in this case take 
up heat from the solution. 

The dissolved substance is called the solute, while the 
liquid in which it dissolves is the solvent. The product 
is a solution. By the concentration of a solution is meant 
the quantity of the solute dissolved in a given volume of 
the solvent; this is usually expressed in grams per liter. 
(See Metric Tables.) 

Everyone knows that not all substances are equally 
soluble, but even the most soluble reaches a limit when 
no more of the substance will go into solution without 
heating. "When this condition is reached. the solution is 
said to be saturated. Anything short of this concentra- 
tion is called an unsaturated solution. A hot saturated 
solution is sometimes cooled without depositing any of its 
solute. A solution in this condition is said to be super- 
saturated. It is easy to tell which of these conditions a 
solution is in, since an unsaturated solution will dissolve 
more of the solute, a saturated one will not dissolve more 
without heating, while a supersaturated solution on the 
addition of more of the solute will deposit its excess and 
become a saturated solution. Heat is evolved during 
this process. 

Water is not the only liquid which will dissolve sub- 
stances ; many others, as alcohol and kerosene, will do the 
same thing. Water, however, is the best solvent known; 
it will dissolve the greatest number of substances. Few 
things are entirely insoluble in water, though many sub- 
stances are nearly so. 

Drinking" Water. — It is this great solvent power of 
water which is responsible for much of the impurity of 
drinking water. Rain falls upon the earth and, sinking 
down into the ground, dissolves many things which ren- 
der it impure. But there is another trouble with drink- 
ing water ; all over the ground everywhere, and often in 



WATER 



9 



the air, are innumerable little organisms so small that 
they can be seen only with a good microscope. Some of 
these bacteria are the causes of disease, others cause decay 
in food and other things, some produce the souring of 
milk, fruit, juices, etc. These little plants grow and mul- 
tiply very fast. This is particularly true where animal 
matter is decaying. You see 
from this why it is so danger- 
ous for drinking water to be- 
come contaminated by the 
drainage from stables, sew- 
ers, and the like, and that 
impure water is often the 
cause of sickness. All drink- 
ing water, then, unless known 
to be pure, should be puri- 
fied. 

Purification of Drinking 
Water. — The simplest way 
to do this is to boil the water, 
since boiling kills the bac- 
teria, but this is not the most 
desirable method, for boiled 
water does not have a pleas- 
ant taste. It must also be 
remembered that boiling will 
not remove all of the mineral 
matter dissolved in the water, 
though sometimes it will re- 
move part of it. A process 

which will remove both the bacteria and mineral matter, 
and which is known as distillation, consists of changing 
water into steam and then condensing the steam into 
water again by passing it through long coils of cold pipe. 
The pipe is kept cold by surrounding it with cold water. 




=jLO Faucet-Water 
^rO 3 inlet 



Distilled Water 
Outlet 

Fig. 4. — Sectional. Diagram of a 
Still, fob Distilling Water in 
Small Quantities. Water in a 
boiler is heated by a gas burner and 
converted into steam. This steam 
passes out through a long pipe 
which is cooled by the faucet-water 
in the surrounding jacket. The 
faucet-water thus becomes partly 
heated before entering the boiler. 
Since it requires more water to con- 
dense the steam than is required for 
the boiler, the excess flows away 
through an outlet pipe. 



10 



A PRACTICAL CHEMISTRY 



Distillation will not remove gases or volatile substances, 
since they will pass over with the steam. 

The water supply of large cities is often purified by 
letting it sink slowly through large filters. The filter 
forms the bottom of a reservoir. The water in the reser- 
voir soaks down through a bed of sand, or of sand and 
charcoal, then through gravel and broken stone, and col- 




Fig. 5. — Explanation of Sand Filter. R, river from which the water supply comes 
through the inlet pipe, I. SR, settling reservoir in which the water is allowed to 
stand till the sediment has fallen to the bottom. The water next passes into the 
filtering reservoir 3 FR, where it sinks slowly through the thick layers of sand, S, 
and gravel, G, and collects among the large rocks, B. The outlet pipe, O, conducts 
the water into the distributing reservoir, DR, from which it flows into the city mains. 



lects in the open spaces between large stones at the bot- 
tom. Then it drains away into a reservoir below, from 
which it flows into the waterpipes of the city. Of course, 
the sand soon becomes dirty and must often be replaced 
by a fresh supply. 

Crystals. — If you examine the "rock candy' ' sold in 
the stores, you will find it made up of little pieces clus- 
tered together. These little bits are all much alike in 



WATER 



11 



shape, they all have the same number of sides, and these 
sides are bright and shining. These bits of candy are 
sugar crystals. They were formed in a sugar solution by 
the molecules of sugar coming together as the water 
evaporated into the air. We may say that crystals are 
regular geometrical solids formed by natural processes. 




Fig. 6. — Sand Filter at Washington, D. C, Before the Sand Was Put In. Loosely 
jointed pipes instead of large stones are placed at the bottom of this filter. The 
water after passing through the sand enters these pipes and flows into a storage 
reservoir. The filter is covered by an arched roof level with the ground and serves 
as a small park. 



Among these natural processes may be mentioned first the 
evaporation of a solution as illustrated by the sugar solu- 
tion given above. Again it has been intimated that solu- 
tions on cooling often deposit crystals. Many solids, when 
melted and allowed to cool slowly, solidify as a mass of 
crystals. Illustrations of this are found in many rocks, 
quartz being an excellent example. Another way in which 
crystals are formed is by the sudden cooling of a vapor. 



12 A PRACTICAL CHEMISTRY 

Snow is an illustration of this ; the water vapor in the 
upper air crystallizes before it has time to form raindrops. 

Water of Crystallization. — Sometimes, when crystals 
form, one or more molecules of the water of the solution 
combine with each molecule of the dissolved substance 
to form the crystal. "Water thus combined in a crystal is 
called water of crystallization, and has lost the properties 
of ordinary water. The crystals are perfectly dry to the 
touch. If the crystals are heated, however, the water 
will escape as steam. Some crystals containing water 
of crystallization lose this water without heating if they 
are exposed to the air. As the water escapes the crystal 
falls into a powder. Crystals which thus lose their water 
are said to be efflorescent and the process is called efflores- 
cence. Many kinds of crystals do not contain water of 
crystallization. 

Many substances absorb water from the air in consid- 
erable quantity and dissolve in it; such substances are 
called deliquescent. 

Composition of Water. — If we are to follow the plan 
of the small boy with a new toy, as mentioned above, 
having now examined water to some extent and found 
out a little of what it will do, we shall next try to break 
it up. To do this we shall need the help of an electric 
current and some sulphuric acid, also the apparatus shown 
in Figure 7. The tubes, H and 0, are united at the bottom 
with the reservoir tube, R, and through R are filled with 
a mixture of water and sulphuric acid in about the pro- 
portion of one part of acid to twenty parts of water. 
Fused through the glass at E and E' are two wires ter- 
minating in platinum plates. These wires lead back to 
a source of electric current, such as a battery or dynamo. 
When the current is turned on, small bubbles of gas form 
on the platinum plates and pass up the tubes. Soon it is 
found that the gas is collecting twice as fast in one tube 



WATER 



13 




as in the other. When a flame is brought near to this gas 
it catches fire and burns with an almost colorless flame. 
The gas in the other tube will not catch fire, but many 
things will burn more freely in it than in the air. In 
many respects the two gases are quite different, but 
neither can be separated into anything simpler; so Ave 
must regard them as elements. The 
one which is in larger volume is 
hydrogen and the other one is called 
oxygen. 

Many other experiments have 
been made to find out the com- 
position of water, and these all 
lead to the conclusion that two 
volumes of hydrogen are combined 
with one of oxygen to form water. 
Oxygen, however, is about sixteen 
times as heavy as hydrogen, so 
that the composition of water by 
weight is about one part of hydro- 
gen to eight parts of oxygen. Be- 
fore proceeding to the further 
study of these gases, which will be 
our next duty, we must stop to 
consider how they exist in water, 
that is, whether they are combined 
or only mixed — whether water is a 
compound or a mixture. 

Compounds and Mixtures. — There are three important 
differences between compounds and mixtures: (1) a com- 
pound always has the same composition — it is always 
made of the same elements in the same proportions; (2) 
the properties of a compound are not the same as those 
of the elements in the compound; (3) you cannot take a 
compound apart by mechanical means. In the case of 




Fig. 7. — The Electroly- 
sis of Water. 



14 A PRACTICAL CHEMISTRY 

mixtures the reverse of these statements is always true. 
Let us try water with reference to these points. "When 
we decomposed water by the electric current we found it 
separated into two gases, and there was twice as large 
a volume of one gas as of the other. This result has been 
verified by many repetitions. This shows that water has 
a definite composition. The properties of the gases are 
very different from those of water. Lastly it requires a 
chemical process to decompose water. It must follow, 
then, that water is a compound, since it answers all these 
three requirements of compounds. Air, on the other hand, 
is a mixture ; its composition is always changing, its prop- 
erties are those of the gases of which it is composed, and 
these gases may be separated by physical means. 



SUMMARY 

General Properties of Water. — Water evaporates, forming water 
vapor which may condense to rain and snow. It freezes at 
0° C, or 32° F., and boils at 100° C, or 212° F. Water is 
at its greatest density at 4° C. and expands as the tempera- 
ture changes in either direction from this point. While 
freezing it expands one-tenth of its volume. Water at 4° C. 
is taken as a standard for determining specific gravity of 
solids and liquids, specific gravity being defined as the 
weight of a given solid or liquid as compared with the 
weight of an equal volume of water at 4° C. 

Solution. — A solution is a liquid in which the molecules of some 
dissolved solid or gas are moving freely. To do this the 
molecules of a solid must take up energy; hence, heat aids 
solution of solids and prevents solution of gases. 

For the same reason pressure aids solution of gases. The 
substance dissolved is called the solute, the liquid is the sol- 
vent, the quantity of solute in a given volume determines 
the concentration of the solution. A solution is saturated 
when its concentration cannot be increased without increase 



WATER 15 

of temperature, unsaturated when its concentration can be 
increased without increase of temperature. A supersaturated 
solution is one which will deposit crystals on the addition 
of a crystal of the solute. 

Drinking water may contain mineral matter dissolved 
from the ground, organic matter from decaying animal and 
vegetable substances, and. bacteria which are the cause of 
decay and disease. Drinking water is purified by boiling, 
by filtering through sand or charcoal, and by distillation, 
which consists of changing water into steam and then con- 
densing the steam again into water. 

Crystals. — Crystals are regular geometrical solids formed either 
by the cooling or evaporation of a solution, the cooling of 
a molten mass, or by the sudden cooling of a vapor. 

Water of Crystallization. — rWater of crystallization is water that 
has lost its ordinary properties by becoming a part of the 
molecules of a crystal. 

Efflorescence. — Efflorescence is the process of giving off water of 
crystallization when exposed to the air. 

Deliquescence. — Deliquescence is the process of absorbing water 
from the air and dissolving in it. 

Electrolysis. — Electrolysis and other experiments show that water 
is composed of two volumes of hydrogen and one volume 
of oxygen, or about eight parts by weight of oxygen to one 
part of hydrogen. 

Compounds. — Compounds are unlike mixtures in that they have 
definite, constant composition, are not decomposed by me- 
chanical means, and do not have the properties of their 
constituents. 

REVIEW QUESTIONS 

1. Make a summary of familiar facts concerning water 
which are not given in the text. 

2. Why does a cloth dry after being wet? 

3. Why does the ocean never become full to overflowing? 

4. Explain the formation of snow. 

5. Describe the appearance of snow flakes. 



16 A PRACTICAL CHEMISTRY 

6. Why does freezing often cause water pipes to leak? 

7. Which is larger, and how much, a pound of water at 
0° C. or a pound of ice at 0° C? 

8. Explain what changes occur in the volume of water dur- 
ing cooling. 

9. At what temperature does water reach its maximum 
density? 

10. Compare the Centigrade and Fahrenheit thermometers. 

11. Calculate at what temperature the reading of these ther- 
mometers is the same. 

12. What is meant by the expression, "The specific gravity of 
iron is 7"? 

13. What becomes of sugar when put into hot coffee? 

14. What is meant by molecules? 

15. How can you show that the molecules of a dissolved sub- 
stance move through the liquid? 

16. Will water dissolve more air in summer or in winter? 

17. Why has the concentration of the Dead Sea increased 
with time? 

18. When a solution of camphor in alcohol is poured into 
water solid camphor is formed. Explain why. 

19. How could this camphor be again brought into solution? 

20. Why does soda water effervesce when drawn into a glass? 

21. Why does it often cool water to dissolve a substance in it ? 

22. How could you tell whether sulphur is soluble in water? 

23. Why is drinking water apt to be impure? 

24. Name the most dangerous impurities that often occur in 
drinking water. 

25. Compare the various methods of purifying drinking water, 
and state under what conditions each is desirable. 

26. How would you make crystals of a metal? 

27. What is water of crystallization? 

28. What kind of substances make good drying agents? 

29. Describe one experiment by which the composition of 
water may be determined. 

30. Explain how you can tell whether candy is a compound 
or a mixture. 



CHAPTER III 
OXYGEN 

Discovery. — About the time of our American Revolution 
Joseph Priestley, an English chemist, found that on heat- 
ing a red powder, which had been made by heating mer- 
cury in the air, he could obtain a gas. If you will heat 
some of the same red powder, which we now call mercuric 
oxide, you may obtain the same gas which Priestley dis- 
covered. Priestley did not name his new gas oxygen, but 
it received this name a little later from the French chem- 
ist, Lavoisier. The discovery of oxygen is an important 
event in the history of chemistry, since it led to the true 
explanation of combustion and to the overthrow of the 
false theory held at that time. 

Occurrence. — Since the red powder from which Priest- 
ley obtained oxygen was made by heating mercury in the 
air, the oxygen which he obtained must have come orig- 
inally from the air. Hence, air contains oxygen. We have 
seen (see p. 13) that it is an important part of water. 
It is also a constituent of nearly everything that grows. 
In the solid earth no other element is found in such large 
quantities or in so many compounds. You see from this 
that it is a very important substance. 

Properties. — From our knowledge of the air we know 
some of the properties of oxygen. It is colorless and has 
no taste or odor. In the air it is mixed with other gases 
lighter than itself, so that air is not quite so heavy as 
oxygen. One liter of oxygen at 0° C. and normal at- 

17 



18 A PRACTICAL CHEMISTRY 

mospheric pressure weighs 1.43 grams. If you will thrust 
a splinter with a spark on it into a jar of oxygen, the 
splinter will at once burst into flame. From this you see 
that oxygen makes substances burn. It does this by unit- 
ing with the elements of the burning substance, producing 
new compounds called oxides. 

Combustion. — The oxygen of the air acts in the same 
way, but the other substances in the air get in the way 
and thus make the burning less rapid than in pure oxygen. 
A fire burns faster when supplied with plenty of air, but 
will go out if the air is entirely shut off. The drafts of 
our stoves and furnaces regulate the flow of air through 
the fire. The fuel (the coal or wood) unites with the 
oxygen of the air and produces, besides the ash, certain 
gaseous oxides which pass up the chimney. If there is 
not enough air, some of the fuel may also pass up the 
chimney. This partly explains the great clouds of smoke 
given off in burning soft coal. There is evidently much 
loss when a fire is giving off much smoke, but fuel may 
also go up the flue in the form of invisible unburned gases. 
Most fires require some heat to start them ; the substances 
must be brought to the kindling temperature before they 
will burn. This is usually done with a match flame. The 
kindling temperature is not the same for all substances ; 
thus wood burns at a much lower temperature than iron, 
and phosphorus at a lower temperature than wood. 
Again, if fuel is cut into small pieces it will burn more 
freely. This is because it presents more surface- to the 
oxygen of the air, and requires less heat to bring the 
pieces to the kindling temperature. A few substances 
burn spontaneously, that is, they start to burn without 
the aid of man. Many cases of so-called spontaneous 
combustion may, no doubt, be traced to other causes. If 
a substance burns spontaneously, it is because slow chemi- 
cal action has been going on for some time and the heat 



OXYGEN 



19 



thus produced has accumulated uutil the substance being 
acted upon is heated to the kindling temperature. The 
substance then bursts into flames. Green hay stored in 
barns is said sometimes to take fire in this way. 

Fire suggests to us two things — light and heat, so we 
may say that burning, or combustion, is chemical action 




Fig. 8. — Generating Oxygen. In the test tube, T, the mixture of potassium chlorate 
and manganese dioxide is heated. The oxygen passes through the rubber delivery 
tube displacing the water in the bottle, O, over the pneumatic trough, P. 



producing light and heat. As we have already tried to 
show, ordinary combustion is chemical action in which the 
burning substance is uniting with oxygen, that is, it is 
very rapid oxidation. In your experiments with oxygen 
you may find that even metals burn in pure oxygen and 
form oxides. These oxides are not so very different from 
the rust you find on metals exposed to the air. In the 
air, however, moisture and some of the other gases present 
help to cause the rusting. As we study the various ele- 



20 A PRACTICAL CHEMISTRY 

ments we shall find many different oxides and different 
kinds of oxidation. 

Preparation. — Priestley obtained oxygen from mercuric 
oxide. We have seen that it may be liberated from 
water; there are besides these many other substances 
from which it may be obtained. "We shall, however, give 
our time to but one method. For laboratory use oxygen 
is prepared by heating in a tube a mixture of three parts 
of potassium chlorate and one of manganese dioxide. 
The gas is collected in bottles over water as shown in 
Figure 8. 

Potassium chlorate is a compound of. the elements, 
potassium, chlorine and oxygen. When heated it decom- 
poses, the oxygen passing off as a gas and leaving the 
potassium and chlorine united as a compound called po- 
tassium chloride. (Notice the difference in the endings 
of the two words chlorate and chloride.) The manganese 
dioxide remains unchanged, but its presence causes the 
potassium chlorate to decompose at a much lower tem- 
perature than when heated alone. 

Oxygen and Life. — Oxygen is necessary for life, since 
breathing is taking oxygen into the lungs, where it passes 
through the tissue into the blood. Here it combines with 
the elements of food and with waste products, furnishing 
energy to the body. In cases of sickness, where the lungs 
are partly filled up and cannot receive enough air to sup- 
port life, pure oxygen is supplied to the patient. For this 
purpose oxygen may be purchased in large metal 
cylinders. 

In the lungs oxygen often serves another purpose be- 
sides those mentioned above. While one is breathing im- 
pure air many kinds of bacteria enter the lungs, and some 
of these, especially those of tuberculosis, find lodgment. 
By breathing large quantities of pure air these bacteria 
may be destroyed. Hence the importance of keeping the 



OXYGEN 



21 



head erect and the shoulders well back that all parts of 
the lungs may be often filled with fresh air. 

Decay. — "We are all familiar with the spoiling or decay 
of a]l kinds of food and of many other things. This is 
brought about by the growth of bacteria aided by the 
oxygen of the air. Moisture is also necessary. The sub- 
stance in which the decay takes place is decomposed into 




Fig. 9. — Generating a Tank of Oxygen. The potassium chlorate and manganese 
dioxide are heated in the copper retort. The tank is first filled with water; as the 
oxygen enters it forces the water out. 



other compounds, some of which may have bad odors. 
These compounds either escape into the air or are carried 
by water into the earth. 

It is natural to think of decay as the enemy of life and 
growth, but this is not the case. The earth would soon 
become covered with useless material if it were not re- 
moved by decay, while the substances thus oxidized and 
decomposed are converted into compounds which are food 
for plants. In like manner the waters of rivers and lakes 



22 



A PRACTICAL CHEMISTRY 



and of the ocean are purified by the oxygen of the air. 
The continuous motion of the water as the wind blows it 
into waves or as it runs down hill brings the organic 
matter which it contains into contact with the oxygen. 
Hence the purity of running water as compared with 
ponds. 

Ozone. — It may seem a little strange that an element 
may exist in more than one form, but this is true of sev- 




Fig. 10. — Making Ozone. The water in the upper tank forces the oxygen from the 
lower tank through the drying bottle containing sulphuric acid and into the long 
glass tube. Within this tube is a coil of aluminium wire, while around the outside 
of the tube is another coil of aluminium wire. A silent discharge of electricity 
from the induction coil, to which the two coils of aluminium wire are connected, 
changes the oxygen into ozone. The ozone escapes from the end of the glass tube. 



eral of the elements, and oxygen is one of this number. 
The other, or peculiar, form of oxygen is called ozone. 
Its properties are quite different from those of ordinary 
oxygen. It has a strong, irritating odor, is one and one- 
half times as heavy as oxygen, is easily decomposed, and 
is much more active in combining with other substances. 



OXYGEN 23 

Many kinds of bacteria are killed by ozone. The mole- 
cules of ozone are larger and heavier than those of oxy- 
gen. It can be easily formed by discharging electricity 
silently through dry air. This is illustrated in Figure 10, 
in which the oxygen stored in the large gas holder is 
dried by passing through sulphuric acid in the wash bottle 
and then goes into the glass tube. In this tube is a coil 
of wire joined to one binding post of the induction coil, 
while around the tube is another coil which is connected 
with the other post of the induction coil. The electricity 
in passing from one coil to the other changes the oxygen 
into ozone. The ozone escapes at the stopcock at the end 
of the glass tube. 

Ozone is formed in small quantities in many oxidation 
processes, and no doubt is often in the air in small quan- 
tities. 

SUMMARY 

Oxygen was discovered by Priestley in 1774 and was named by 
Lavoisier in 1778. It is an important constituent of water, 
oxides, and other compounds. It is part of many rocks, 
of nearly everything that grows. Uncombined oxygen occurs 
in the air mixed with other gases. Oxygen is colorless, 
tasteless and odorless. One liter of it under standard con- 
ditions weighs 1.43 grams. 

Combustion. — Combustion is any chemical action producing light 
and heat. The chemical action of most combustion is oxida- 
tion. 

Kindling Temperature. — Kindling temperature is the temperature 
to which fuel must be brought before burning begins. 

Spontaneous Combustion. — Spontaneous combustion is combus- 
tion produced by chemical action without the aid of man. 

Oxygen and Life. ^-Oxygen may be generated by electrolysis of 
water, and by the decomposition of oxides and of other oxy- 
gen compounds. For laboratory use it is best obtained by 
heating potassium chlorate. If manganese dioxide is mixed 



24 A PRACTICAL CHEMISTRY 

with the potassium chlorate, the oxygen is liberated at a 
lower temperature. 

The oxygen of the air is necessary for life. Pure oxygen 
is administered to those not able to obtain it in sufficient 
quantity from the air. 

Decay. — Decay is necessary in the economy of nature. Oxygen 
is an important factor in decay and in the natural purifica- 
tion of the waters of streams, lakes, and oceans. 

Ozone. — Ozone is a peculiar form of oxygen. It has more energy 
and is much heavier than ordinary oxygen. It may be 
formed by a silent discharge of electricity through dry 
oxygen. 

REVIEW QUESTIONS 

1. When was oxygen discovered? 

2. How did its discovery affect the theories of combustion? 

3. If you were naming this gas, what would you call it? 

4. How can you show that oxygen is in the air? 

5. What can you say of the quantity of oxygen in the earth ? 

6. Name five properties of oxygen. 

7. What are oxides? 

8. Why does a fire go out when the air is shut off? 

9. What conditions cause loss of fuel. 

10. Why do not most substances burn spontaneously? 

11. Under what conditions do substances burn spontaneously? 

12. How does combustion differ from oxidation? 

13. Are combustion and oxidation ever the same thing? 

14. How may oxygen be prepared for laboratory use? 

15. Why are the jars in which oxygen is collected first filled 
with water? 

16. Why is manganese dioxide used in this process? 

17. What advantages and disadvantages would result if the 
air were pure oxygen? 

18. What is meant by decay? 

19. What are bacteria? 

20. Why is decay a good thing? 

21. How does ozone differ from ordinary oxygen? 

22. How can it be made? 



CHAPTER IV 
HYDROGEN 




Recalling the experiment with water and the electric 
current, you will remember that two gases were obtained. 
One of these, oxygen, has been the subject of our last 
chapter; the other one, hydrogen, we 
shall now consider. 

Properties. — Like oxygen, hydrogen 
is a colorless, odorless, tasteless gas, 
but, unlike oxygen, it will burn (with 
an almost invisible name) and will not 
permit most other substances to burn in 
it. Carefully notice this point — many 
substances burn freely in oxygen, while 
but few burn in hydrogen. It may be 
shown that hydrogen burns freely in 
air, but does not permit a taper to burn 
in it, by lighting a taper and thrusting 
up into an inverted jar of hydrogen as 
in Figure 11. The hydrogen burns at 
the mouth of the jar, but the taper does 
not burn while up in the jar. When 
the taper is removed it lights again 
while passing through the hydrogen 
flame. The jar of hydrogen must be held mouth down be- 
cause air is 14.4 times as heavy as hydrogen, and the hydro- 
gen would escape if the jar were turned up. Oxygen is 

25 



Fig. 11. — Hydrogen 
not a Supporter 
of Combustion. 
When the lighted 
splint, S, is thrust 
up into a jar of hy- 
drogen, the flame on 
the splint is extin- 
guished but the hy- 
drogen catches fire 
and burns in the air 
at the mouth of the 
jar. 



26 



A PRACTICAL CHEMISTRY 



nearly sixteen times as heavy as hydrogen. In fact hydro- 
gen is the lightest gas known. One liter of hydrogen under 
standard conditions weighs 0.08984 gram. It is this prop- 
erty which makes it useful for filling balloons. That hy- 
drogen is so much lighter than air is well illustrated by 
pouring it upward as in Figure 12. 

A jet of oxygen will burn as freely in a jar of hydrogen 
as hydrogen will burn in oxygen or air. This may be 
shown by a simple experiment illustrated in Figure 13. 

Although the hydrogen 
flame gives so little light, it 
is very hot. This is partic- 
ularly true if oxygen is 
mixed with the hydrogen in 
the flames. The apparatus 
for doing this is called the 
oxyhydrogen blowpipe, and 
consists of two tubes, one 
within the other, as shown 
in Figure 14. The outer 
tube carries the hydrogen, 
and the inner tube the oxy- 
gen. The two gases are 
stored in tanks. The hydro- 
gen is lighted first, and then the oxygen is turned on. A 
cylinder of lime, if put into the oxyhydrogen flame, be- 
comes white hot and gives a brilliant light, known as the 
limelight sometimes used in stereopticons. This blowpipe 
flame is also used to melt metals, where a high heat is 
necessary, as, for example, in the welding of lead plates 
in acid chambers, tanks, etc. 

"When hydrogen burns in the air or in oxygen it unites 
with the oxygen in the proportion of about 8 parts by 
weight of oxygen to 1 of hydrogen, forming water. Of 
course this water is in the form of steam and at a very 




Fig. 12. — Pouring Hydrogen. The hy- 
drogen in A is poured upward into B. 
Its presence is afterwards tested by 
bringing the mouths of A and B suc- 
cessively to a flame. 



HYDROGEN 



27 




high temperature. Steam has great expansive power, 
and this power increases very much as the temperature 
of the steam is increased. Hence, 
when steam is formed at the high tem- 
perature of the hydrogen flame it ex- 
pands with great force. If the hydro- 
gen and oxygen (or hydrogen and air) 
are mixed before being lit, the expan- 
sion of the steam is so sudden and vio- 
lent that an explosion is the result. 
Many students have been injured by 
carelessly lighting jets of hydrogen 
mixed with air. This danger may be 
avoided by collecting some of the hy- 
drogen in a test tube, lighting the hy- 
drogen in this test tube and with this 
burning hydrogen lighting the jet of 
hydrogen. 

The fact that water is produced 
when hydrogen burns in the air may 
be shown by means of an apparatus 
arranged as shown in Figure 15. The 
hydrogen from the Kipp apparatus, A, 
is passed through the drying tube, D, 
filled with calcium chloride. The dry 
hydrogen thus obtained is burned in 
the jet, F, and the steam condenses on 
the cold glass held above. 

Reducing Agents.— This strong af- 
finity which hydrogen has for oxygen 
causes it to unite not only with free 
oxygen but also with oxygen which is 
in other compounds. Any substance which will thus unite 
with the oxygen of compounds and remove it is called a 
reducing agent. Hydrogen is a good reducing agent. 



Fig. 13. — Burning Air 
in Illuminating Gas. 
The glass bulb is first 
filled with illuminating 
gas and lighted where 
it escapes at the top. 
As the gas thus passes 
through the bulb it 
draws in a current of 
air through the short 
tube. This stream of 
air may be lighted by 
putting a small flame 
up through the short 
tube. We thus have 
an air flame burning 
in an atmosphere of 
illuminating gas. A 
similar result might be 
obtained by using hy- 
drogen in place of 
illuminating gas and 
oxygen in place of air 




28 A PRACTICAL CHEMISTRY 

This may be illustrated by heating some copper oxide in a 
glass tube and at the same time passing dry hydrogen 
through the tube. The hydrogen will take oxygen from 
the hot copper oxide and form water. 

Preparation. — ¥e have learned (see p. 13) that hy- 
drogen may be obtained from water by the aid of the 
electric current. There are other ways of decomposing 
water and driving out part or all of the hydrogen. If 
steam is passed through a hot iron tube, oxygen of the 
steam unites with the iron to form an oxide of iron, while 

, the hydrogen escapes 

— *^? ^ ^^ and may be collected 

over water in a jar. The 
same thing is true when 
some of the other metals 

Fig. 14. — The Oxy-hydrogen Blowpipe. 

This consists of two tubes, one within are used at high tempera- 

the other. The inner tube through which ^^ There &re algQ 

the oxygen passes extends nearly to the 

end of the outer tube. Let the student SOme metals which will 

think out all the reasons why the oxygen ^ ecompose water violent- 

should enter the flame from the inside m x 

tube rather than from the outside tube. ly without being heated. 

Sodium is one of these. 
If a small piece of sodium be thrown upon a dish of water, 
it will swim about with a sputtering sound. If we wrap the 
sodium in a bit of paper and put it under the mouth of 
an inverted test tube of water in a dish of water, we shall 
soon see bubbles rising in the tube. These bubbles are 
hydrogen which the sodium has driven out of the water. 
Sodium removes one-half of the hydrogen from water 
and forms a compound with the oxygen and other half 
of the hydrogen. This compound is called sodium hy- 
droxide. We shall have occasion to study it a little later. 
The most convenient way to generate hydrogen for 
laboratory purposes is to drive it out of an acid by means 
of a metal. All acids contain hydrogen, and many of the 
metals will drive the hydrogen out of acids and take 




x-jk 



Fig. 15. — Water From Burning Hydrogen. Hydrogen is generated in the Kipp 
apparatus by means of acid from the reservoir, A, flowing down the long tube 
into B and thence up into H where it comes into contact with zinc whenever the 
gas pressure is relieved by opening the stop-cock, S. When this stop-cock is closed 
the pressure of the generated hydrogen again forces the acid back into A. The 
hydrogen is dried by passing through the calcium chloride tube, D. It burns in 
the flame, F, and the steam passes into the bulb, G, where it is condensed to drops 
of water. 






- =j 


^L 


i i 



Fig. 16. — Decomposition of Steam by Heated Iron. Steam is generated by heating 
water in a flask. The steam passes through the iron pipe filled with bits of iron 
and heated by the gas flames below. The oxygen of the steam unites with iron, 
forming an oxide of iron, while the hydrogen is free to pass out into the cylinder 
at the right where it may be collected over water. 



30 A PRACTICAL CHEMISTRY 

its place, but the most satisfactory acids to use are hydro- 
chloric and sulphuric, and zinc is the best metal with 
which to drive out the hydrogen. The apparatus may 
be arranged as shown in Figure 17. The zinc is put into 
the flask, F, and the acid is poured in through the thistle 
tube which reaches nearly to the bottom of the flask. If 
sulphuric acid is used, enough water should first be added 
to cover the zinc, and the acid should be poured in, a 
little at a time, as needed. The hydrogen, when found 
to be free from air by lighting as described above, may 
be collected in bottles over water as in the case of oxygen. 
In the Kipp apparatus, shown in Figure 15, hydrogen is 
generated in the same way as we have just described, but 
the apparatus is automatic in its action. The zinc is put 
into the middle bulb, while the acid is poured in through 
the top and passes down by the inner tube into the bottom 
bulb and then up into the middle bulb, where it reacts 
with zinc, liberating hydrogen. "When the hydrogen pres- 
sure overcomes the pressure of the liquid in the tube the 
acid is forced down away from the zinc, and the action 
ceases. 

While zinc and hydrochloric acid are perhaps the most 

convenient materials to use for generating hydrogen, some 

other metals and acids may be used. Among these may 

be mentioned tin and hydrochloric acid, iron and dilute 

sulphuric acid, and iron and acetic acid. Such metals as 

lead and copper cannot be used, nor can we use nitric acid. 

Hydrogen Peroxide. — Water is an oxide of hydrogen 

containing hydrogen and oxygen in the proportion of 

about eight parts oxygen to one part hydrogen. The 

chemical name for water is hydrogen monoxide. There is 

another oxide of hydrogen, known as hydrogen dioxide, 

in which there is twice as much oxygen in proportion to 

the hydrogen. It has received the name hydrogen dioxide 

to show that it contains twice as much oxygen, since "di" 



HYDROGEN 



31 



means "two." It is also called hydrogen peroxide. Hy- 
drogen peroxide is a strong oxidizing agent, that is, it 
gives up oxygen freely to oxidize other substances. In 
this respect it is quite different from water. In fact, in 
many respects it is unlike water, and only resembles water 
in composition and in being a liquid. Hydrogen peroxide 
has a disagreeable taste, quite noticeable when used in 
dilute solution as a mouth wash or as a gargle for sore 
throat. It is much used in medicine to wash sores. Its 




Fig. 17. — Generating and Collecting Hydrogen. The flask, F, contains the zinc, 
water and hydrochloric acid. The thistle tube, T, should extend nearly to the 
bottom of the flask while the delivery tube, D, should not be more than half an inch 
through the cork. R is a rest to support the gas bottle while being filled, in lieu 
of this the bottle may be held by an iron ring or by the hand. 



cleansing power depends upon the ease with which it de- 
composes and gives oxygen to other things. The solution 
of hydrogen dioxide, as ordinarily sold, contains only 
about three per cent, of the dioxide. The rest is water 
and a very small amount of other substances. More con- 
centrated solutions are very powerful and would seriously 
injure the flesh if put upon it. Hydrogen peroxide is 
made by the action of acids upon barium dioxide. It can- 
not be made by the direct union of hydrogen and oxygen. 
Importance. — We have studied a few compounds con- 
taining hydrogen; there are many others. Hydrogen is 



32 A PRACTICAL CHEMISTRY 

an important element in everything that grows, whether 
plants or animals. It is a constituent of all acids. It is 
also in many minerals. Of course, all elements are im- 
portant, and each has its own place in nature, but the part 
played by hydrogen is very large. 



SUMMARY 

Hydrogen is a colorless, odorless, tasteless gas which will burn 
freely in air, but will not permit many things to burn in it. 
Oxygen, however, burns freely in hydrogen. Mixtures of 
these gases burn with explosive violence, due to the expan- 
sion of the highly heated steam formed by the combustion. 
The hydrogen flame is very hot, but gives little light. The 
oxyhydrogen blowpipe is used where very high tempera- 
tures are required. Hydrogen is the lightest substance 
known, weighing under standard conditions 0.08984 grams 
per liter. Water is the only product of the combustion of 
hydrogen in air or in oxygen. 

A Reducing Agent. — A reducing agent is a substance which will 
unite with the oxygen of compounds and thus remove it. 
Hydrogen is a good reducing agent. Hydrogen may be gen- 
erated by the electrolysis of water, by the action .of metals 
upon water, and by the action of metals upon acids. Hy- 
drogen forms two oxides — water and hydrogen peroxide. 
The latter contains twice as much oxygen in proportion to 
the hydrogen as there is in water. Hydrogen peroxide is a 
strong oxidizing agent, that is, it gives up oxygen freely to 
oxidize other substances. Hydrogen is an element of much 
importance, since it is a constituent of many things. 



REVIEW QUESTIONS 

1. Of what elements is water composed? 

2. In what respects does hydrogen resemble oxygen? 

3. In what respect does it not resemble oxygen? 



HYDROGEN 33 

4. What method would you use for generating hydrogen in 
the laboratory? 

5. Describe an experiment to show that ordinary fuel will 
not burn in hydrogen. 

6. Why does not an explosion occur when the oxyhydrogen 
flame is lighted? 

7. The sun's atmosphere contains much hydrogen; did ours 
ever contain hydrogen? 

8. How do you explain hydrogen explosions? 

9. How may these explosions often be avoided? 

10. Why is steam given off from most flames? 

11. By what method would you fill a balloon with hydrogen? 

12. How do you collect a jar of hydrogen? 

13. How does hydrogen peroxide differ from hydrogen mon- 
oxide? 

14. Compare oxidizing and reducing agents. 



CHAPTER V 

ACIDS, BASES, AND SALTS 

While studying oxygen a few oxides were mentioned; 
let us remember that nearly all the elements unite with 
oxygen. These oxides have many interesting properties. 
Many of them will dissolve in water and form compounds 
with the water. The compounds thus formed are of two 
very different classes. 

ACIDS 

It would be interesting to make a large number of these 
oxides, dissolve them in water and study the compounds 
formed. If, for example, we were to dissolve the oxides 
of sulphur, nitrogen, phosphorus, and many others, we 
would find that the solutions all have a sour taste and 
will turn blue litmus red. The compounds of this class 
are acids. Besides having a sour taste and turning blue 
litmus red, acids have various other properties ; they cor- 
rode metals, decompose carbonates, are often injurious to 
the flesh and clothing, change many organic dyes, and 
neutralize hydroxides. Some acids, are liquids, others 
are solids, while still others are gases. You are familiar 
with some acids : in vinegar there is acetic acid, lactic acid 
is in sour milk, citric acid is in lemons, while in "soda 
water" we have carbonic acid. While discussing the 
preparation of hydrogen mention was made of a few acids 
and of the fact that they all contained hydrogen. Some 
of them also contain oxygen, but this is not a necessary 

34 



ACIDS, BASES, AND SALTS 35 

part of acids, while hydrogen is a necessary part. Oxy- 
gen means "acid former," but hydrogen is the true acid 
former. 

Naming" of Acids. — If each element formed but one acid, 
the whole matter would be easy; the acid of chlorine 
would be chlorine acid and that of sulphur would be 
called sulphur acid. This, however, is not the case since 
many of the non-metallic elements form one or more acids 
with hydrogen alone, and also several with hydrogen and 
oxygen together. Each of these must have a name, and, 
in order that the names of all acids may be uniform, it 
has been necessary to establish some rules. The names 
of all acids which do not contain oxygen begin with 
hydro- and end in -ic. Thus the acid of chlorine and hy- 
drogen is hydrochloric acid. There are quite a number 
of these acids, as hydrosulphuric, hydrofluoric, hydro- 
bromic, etc. In naming the oxygen acids of an element 
it is customary to give the most common or best known 
one of the series a name ending in -ic ; the acid with more 
oxygen than this one receives the same name with per- 
placed before it. Thus the best known of the oxygen 
acids of chlorine is called chloric acid, and the acid with 
more oxygen than chloric acid is perchloric acid. The 
acid with less oxygen than the -ic acid is given a name 
ending in -ous, and the one with still less oxygen receives 
the same name with the prefix hypo-. As illustrations of 
this we have chlorous and hypochlorous acids. The 
chlorous acid has less oxygen than the chloric acid, and 
the hypochlorous still less than the chlorous. 



BASES 

The other class of compounds formed when oxides are 
dissolved in water is known as bases. They are about 
the opposite of acids in many respects. They turn red 



36 A PRACTICAL CHEMISTRY 

litmus blue, and have a bitter taste. They are slippery 
to the touch. They always contain both hydrogen and 
oxygen. The hydrogen and oxygen are joined together; 
the oxygen is also joined with the other element, which 
is usually a metal. The hydrogen and oxygen combined 
together in these compounds form what is known as the 
hydroxyl group, and from this the bases are called hy- 
droxides. In naming the bases the name of the metal 
is given first and this is followed by "hydroxide." The 
compound formed by dissolving calcium oxide in water 
is named calcium hydroxide, while sodium oxide gives 
sodium hydroxide. This is the same compound that we 
made by the action of sodium on water". Some of these 
hydroxides are much stronger bases than the rest and are 
spoken of as alkalies. Some of them eat the flesh and 
on this account are said to be caustic ; thus we have ' ' caus- 
tic soda" and "caustic potash." 

NEUTRALIZATION AND SALTS 

When an acid and a base are brought together the hy- 
drogen and oxygen of the base unite with the hydrogen 
of the acid and form water. The properties of both acid 
and base are thus destroyed, that is, the acid and the base 
are neutralized and the process is called neutralization. 
The metal which was in the base now takes the place 
which was occupied by the hydrogen in the acid and a 
new compound called a salt is formed. Ordinary table 
salt is a good example of this class of compounds. They 
crystallize from solution, have various colors and taste 
more or less like common salt. "We may define a salt as 
a compound in which a metal has replaced part or all of 
the hydrogen of an acid. 

It is possible to make salts in a number of other ways 
besides the one just mentioned. Only two of these need 



ACIDS, BASES, AND SALTS 37 

be taken up now. (1) An oxide may be used in place of 
a hydroxide. In this case the oxygen of the oxide unites 
with hydrogen of the acid to form water. (2) A metal 
may drive out the hydrogen of an acid. We may illus- 
trate the three methods thus : 

(1) Zinc hydroxide + hydrochloric acid = water 

-j- zinc chloride. 

(2) Zinc oxide -f- hydrochloric acid = water -f- 

zinc chloride. 

(3) Zinc + hydrochloric acid = hydrogen -|- zinc 

chloride. 

It Jias been mentioned that acids turn blue litmus red 
and bases change red litmus to blue ; there are a number 
of other dyes which change color with acids and bases, 
among which may be mentioned methyl orange and 
cochineal. Such substances are spoken of as indicators, 
because they may be used to indicate whether a substance 
is an acid or a base. They are also used to show when 
the right amount of acid or base has been added to make 
the solution exactly neutral. This process is usually car- 
ried out by means of long graduated tubes called burettes. 
(See Fig. 18.) These tubes have stopcocks at the bottom. 
One tube is filled with acid in solution and the other with 
a solution of the base. A certain quantity of the acid is 
drawn out into a beaker, a few drops of indicator added 
and the base then allowed to flow slowly from the other 
burette until the solution in the beaker is neutral as 
shown by the indicator. It is found by this process that 
it always requires the same quantity of acid to neutralize 
a given quantity of base ; for example, it takes 36.5 grams 
of hydrochloric acid to neutralize 40 grams of sodium 
hydroxide. 

The Naming 1 of Salts. — Since many of the metals form 
salts with many of the acids, the number of these salts 



38 



A PRACTICAL CHEMISTRY 



1^1 



A „ 



©-a 



•F 



Qsn(!) 



must be very great. In naming salts both the metal and 

the acid from which the salt is made must be remembered. 
That is, the first name of the salt must 
be the name of the metal, and the second 
name that of the acid with certain 
changes. The rule may be stated thus: 
In naming salts of acids beginning with 
" hydro-" drop this prefix and change 
the "-ic" ending into "-ide." The 
"-ic" ending of all other acids is 
changed to "-ate" and the "-ous" end- 
ing into "-ite." All other prefixes ex- 
cept "hydro-" are retained. Let us il- 
lustrate this : The salts of hydrochloric 
acid are called chlorides, and those of 
hydrobromic acid are bromides ; thus we 
have sodium chloride, potassium bro- 
mide, etc. The calcium salt of sulphuric 
acid is calcium sulphate, and of chloric 
acid is calcium chlorate, but of per- 
chloric acid it is calcium perchlorate. 
Likewise the sodium salt of chlorous 
acid is sodium chlorite, and of hypo- 
chlorous acid it is sodium hypochlorite. 
A little practice will be necessary in or- 
der to become familiar with the naming 
of acids and salts, but the importance of 
learning to name acids and salts cor- 
rectly and of being able to see quickly 
the connection between these compounds 
can scarcely be overestimated. The stu- 
dent should make it his constant prac- 
tice to establish this connection between 

each new acid and its salts when he is first introduced to 

them. 



Fig. 18. — Neutraliza- 
tion. Two burettes, 
A and B, are sup- 
ported in the stand, 
S. A is filled with a 
dilute acid up to O 
and B with base. 
A definite quantity 
of acid is drawn into 
the flask, C, by 
means of the stop- 
cock, K. A few drops 
of indicator are add- 
ed and the solution 
of base is slowly run 
in until the indicator 
shows that the acid 
has been neutral- 
ized. The flask, C, 
must be shaken with 
a whirling motion 
after each addition 
of base. The vol- 
ume of base used is 
read from the gradu- 
ations of the burette. 



ACIDS, BASES, AND SALTS 39 



SUMMARY 

M«>*t elements form oxides. Oxides, when dissolved in water, 
form acids and bases. Acids are sour, turn blue litmus red, 
decompose carbonates, corrode metals, are often injurious 
to the flesh and clothing. All acids contain hydrogen. 

STEM 

Naming Acids. — prefix ending illustration 

Acids without oxygen . . Hydro- -ic Hydrochloric acid 
Best known oxygen acid -ic Chloric 

With more oxygen than 

in -ic acid Per- -ic Perchloric 

With less oxygen than in 

-ic acid -ous Chlorous " 

With less oxygen than in 

-ous acid Hypo- -ous Hypochlorous " 

Bases. — Bases are almost the opposite of acids. They turn red 
litmus blue and are slippery to the touch. They all contain 
the hydroxyl group. Some bases are caustic. Bases are 
named by adding hydroxide to the name of the metal. 
Neutralization. — Neutralization is the process of destroying the 
characteristic properties of an acid or a base. This is ac- 
complished when an acid and a base are brought together 
by the hydrogen of the acid forming water with the hydroxyl 
of the base. 
Salt. — A salt is a compound in which a metal has replaced part, 

or all, of the hydrogen of an acid. 
Indicators — Indicators are compounds which change color with 

acids and bases. 
Naming of Salts. — Give first the name of the metal, and then 
the name of the acid, with the following changes: 

Hydro- acids. — Drop "hydro-" and change "-ic" to "-ide." 
Other -ic acids. — Change "-ic" to "-ate." 
Acids in -ous. — Change "-ous" to "-ite." 
Retain all prefixes except "hydro-". 



40 A PRACTICAL CHEMISTRY 



REVIEW QUESTIONS 

1. Can all acids and bases be made by the method of mak- 
ing acid and bases given in this chapter? 

2. Name several important characteristic properties of acids. 

3. Name some common acids. 

4. Why cannot hydrogen be obtained from all acids? 

5. Why was oxygen wrongly named? 

6. How are acids which do not contain oxygen named? 

7. Which one of the oxygen acids of sulphur contains the 
least oxygen? 

8. Name the most common oxygen acid containing nitrogen. 

9. Lime water is a base; give its name and peculiar prop- 
erties. 

10. What is the composition of caustic soda? 

11. Name the base formed from magnesium. 

12. What are alkalies? 

13. Describe fully what happens when an acid is neutralized 
by a base. 

14. Define a salt. 

15. Name the silver salt of formic acid. 

16. What is sodium bromide? 

17. What salt is formed when calcium oxide is dissolved in 
nitric acid? 

18. What is the composition of barium carbonate? 

19. What are indicators? 

20. How does the law of definite proportions apply to neu- 
tralization ? 

21. Name the lead salt of hypochlorous acid. 

22. Name the silver salt of hydrofluoric acid. 

23. What elements in sodium sulphide? 

24. What elements in perchromic acid? 



■ 



CHAPTER VI 
, ATOMS, SYMBOLS, FORMULA, AND EQUATIONS 

Atoms. — The composition of any compound, as has been 
stated, is constant. Every sample of pure water has one 
part by weight of hydrogen combined with 7.94 parts of 
oxygen. Likewise in hydrogen peroxide there is one part 
of hydrogen combined with 15.88 parts of oxygen. The 
proportion of oxygen to hydrogen in this second oxide is 
double that in the first. This is no accident ; analysis has 
shown the same relation to exist between the quantities 
of oxygen in the two oxides of carbon, while the quanti- 
ties of oxygen in the five oxides of nitrogen are all simple 
multiples of the quantity in the monoxide. 

A careful consideration of such facts as these, together 
with the fact that matter can be neither created nor de- 
stroyed by chemical action, led the English chemist, John 
Dalton, in 1803 to propound his Atomic Theory. He be- 
lieved: (1) that all elements aire made up of minute in- 
divisible particles or atoms; (2) the atoms of different 
elements have different weights, but all of the atoms of 
the same element have the same weight; (3) all chemical 
changes are either the union or separation of definite num- 
bers of atoms of elements; (4) in chemical changes no 
change can take place in the weight of the atoms. 

This theory explains the definite composition of com- 
pounds, since it requires a definite number of atoms 
having a definite weight. For example, in each molecule 

41 



42 A PRACTICAL CHEMISTRY 

of water there are two atoms of hydrogen of equal weight 
united with one atom of oxygen, also having definite 
weight. 

In like manner it explains why the quantities of ele- 
ments in various compounds are simple multiples of the 
quantities in the simplest compound, as in the case of the 
two oxides of hydrogen and of the various oxides of 
nitrogen mentioned above. Thus hydrogen peroxide has 
two hydrogen atoms combined with two oxygen atoms, so 
that the quantity of oxygen in it is a simple multiple of 
that in water. 

Likewise, in the five oxides of nitrogen we have atoms 
combined in the following proportions: Two atoms of 
nitrogen to one atom of oxygen, two atoms of nitrogen to 
two of oxygen, two atoms of nitrogen to three of oxygen, 
two atoms of nitrogen to four of oxygen, and two atoms 
of nitrogen to five of oxygen. Here again the quantities 
of oxygen are simple multiples of the quantity in the first 
or simplest oxide. 

Since this theory was introduced, more than a hundred 
years ago, many discoveries have been made confirming 
all parts of it except the idea of the indivisibility of the 
atoms. Recent investigations seem to show that some 
atoms at least can be decomposed. If this should be 
proved true of all atoms, it would not destroy the theory 
or impair its usefulness in explaining chemical action. 

Symbols. — It is found convenient in chemistry to use a 
system of shorthand to represent elements, compounds 
and chemical changes or reactions. The fundamental 
characters in this chemical shorthand are called symbols. 
Each symbol represents one atom of an element. The 
symbol is the first letter of the name of the element, ex- 
cept in those cases where more than one element begins 
with the same letter. For these elements the first letter 
and also one of the other letters of the name are used. 



ATOMS, SYMBOLS, FORMULiE, AND EQUATIONS 43 

Sometimes these letters are taken from the Latin name 
of the element. 



H represents one atom of Hydrogen. 





i a 


i i t 


' Oxygen. 


N 


i it 


( t 


1 Nitrogen. 


S 


i it 


a t 


' Sulphur. 


C 


t a 


i t 


1 Carbon. 


CI 


t it 


t t 


' Chlorine. 


p 


t a 


c i 


' Phosphorus 


Mn 


t it 


i i 


1 Manganese. 


K 


i it 


t t 


c Potassium. 


Na 


t it 


t t 


1 Sodium. 


Ag ' 


t a 


t t 


' Silver. 



K, Na, and Ag are derived, respectively, from kalium, 
natrium and argentum, the Latin names of potassium, 
sodium, and silver. 

Formulae. — Molecules are represented by formulas com- 
posed of symbols and figures. Small figures are placed 
between and a little below the symbols to show how 
many atoms of each kind are in the molecule ; for exam- 
ple, H 2 represents one molecule of water containing two 
hydrogen atoms and one atom of oxygen. To indicate a 
number of molecules a large figure is placed in front of 
the formula. 3H 2 S0 4 means three molecules of sulphuric 
acid. Molecules of elements are represented in the same 
way : H 2 is a molecule of hydrogen containing two atoms, 
but 2H means two hydrogen atoms not combined to form 
a molecule. A gaseous element is in this condition of 
single atoms when first liberated from a compound. The 
name to express this condition is nascent state. Ele- 
ments in this condition are very active, reacting with 
many things which they would not attack when in the 
form of molecules. 



44 A PRACTICAL CHEMISTRY 

Equations. — According to the atomic theory, chemical 
action is the union or separation of atoms. These chemi- 
cal reactions are expressed by equations; the formulas 
of the various compounds entering into the chemical 
change are arranged with plus signs between them on 
the left of the equals sign, while those representing the 
products formed are similarly placed on the right. Thus 
the reaction between hydrochloric acid and sodium hy- 
droxide is represented by the equation : 

HC1 + NaOH = H 2 + NaCl. 

This is read, "One molecule of hydrochloric acid reacts 
with one molecule of sodium hydroxide, giving one mole- 
cule of water and one molecule of sodium chloride." In 
like manner the reactions by which ammonia is liberated 
from ammonium chloride are written thus: 

2NH 4 C1 + Ca(OH) 2 = 2NH 4 OH^+ CaCl 2 
and NH 4 OH = NH 3 + H 2 0. 

The first is read, "Two molecules of ammonium chloride 
and one molecule of calcium hydroxide react to form 
two molecules of ammonium hydroxide and one molecule 
of calcium chloride." The second one, "One molecule 
of ammonium hydroxide decomposes, giving one mole- 
cule of ammonia and one molecule of water." 
In the formation of water the equation is, 

2H 2 + 2 = 2H 2 0. 

It will be noticed that in this reaction molecules of 
elements are involved. On the other hand, when hy- 
drogen is liberated, since the gas is in the nascent state, 
it should be represented as separate atoms: 



ATOMS, SYMBOLS, FORMULA, AND EQUATIONS 45 

Zn + H 2 S0 4 = ZnS0 4 + 2H. 
Zinc + Sulphuric Acid = Zinc Sulphate + Hydrogen. 

The student must be warned against all attempts to 
manufacture equations. He must remember that they are 
but a method of expressing facts which have been found 
out by laboratory experiments, and that equations based 
on anything short of experiment are merely guesses. 
Beginners are often led astray by the fact that an equa- 
tion (which they have written) " balances, " that is, they 
find the same atoms in the same numbers on both sides 
of the equation. This is not the least guaranty that the 
equation is correct, for, while a true equation must bal- 
ance, a false one may also do so. 

Valence. — The atoms of all elements do not have the 
same capacity for combining with atoms of other ele- 
ments. An examination of the above formulae will show 
that, while one sodium atom unites with one chlorine 
atom, one atom of calcium unites with two of chlorine. 
One atom of chlorine unites with one atom of hydrogen, 
but one of nitrogen requires three of hydrogen. This 
capacity of atoms to unite with other atoms is called 
valence. 

Hydrogen is taken as the standard by which valence 
is measured. If an atom of an element unites with one 
hydrogen atom, the valence of that element is said to 
be one, or it is a univalent element. If it combines with 
two hydrogen atoms, its valence is two, it is a bivalent 
element. If with three, its valence is three, and it is a 
trivalent element; and so on. 

Those elements that do not combine with hydrogen 
usually replace it, as zinc does in sulphuric acid in the 
above equation. It will be noticed that one zinc atom 
takes the place of two hydrogen atoms, hence the valence 
of zinc is two. "When sodium is put into water one atom 



46 A PRACTICAL CHEMISTRY 

of sodium replaces one atom of hydrogen, as in the 
equation : 

Na + H 2 = NaOH + H, 

hence, the valence of sodium is one. 

We may say, then, that the valence of an element is 
measured by the number of atoms of hydrogen which one 
atom of the element will unite with or replace. 

A number of the elements have two valences and some 
have several. Thus, phosphorus unites with oxygen to 
form phosphorus trioxide, P 2 3 , in which the phosphorus 
has a valence of three and phosphorus .pentoxide, P 2 5 , 
in which the valence of the phosphorus is five. Again 
we have FeCl 2 and FeCl 3 , giving Fe the valences two and 
three, since in the first the Fe replaces the two hydrogen 
atoms of 2HC1 and in the second the three hydrogen 
atoms of 3HC1. FeCl 2 is called ferrous chloride and 
FeCl 3 ferric chloride. It will be noticed that the -ous 
ending is given to the name when the element has the 
lower valence and the -ic ending when it has the higher. 

An element in the lower valence condition may be 
changed to the higher by oxidation, and conversely the 
change from the higher to the lower valence is accom- 
plished by reduction. For example, iron in ferrous chlo- 
ride may be oxidized in the presence of hydrochloric acid, 
or of chlorine, and converted into ferric chloride. On 
the other hand, nascent hydrogen will reduce ferric 
chloride to ferrous chloride. 

It will be seen that, while valence for many elements 
is not constant, its variations are dependent upon definite 
known conditions. 

In writing formulae valence must always be kept in 
mind. 



ATOMS, SYMBOLS, FORMULAE, AND EQUATIONS 47 



SUMMARY 

Elements are made up of minute, indivisible atoms. All atoms 
of the same element have the same weight, but atoms of dif- 
ferent elements have different weights. Chemical action is 
the union or separation of definite numbers of atoms; no 
change can take place in the weight of the atoms. 

Symbols. — Symbols are letters from the name of an element used 
to represent an atom of that element. 

Formulae. — Formulae are combinations of symbols and figures 
used to represent molecules. 

The Nascent State. — The nascent state is the condition of a 
gaseous element when first liberated from a compound. It 
is then composed of single atoms which have not had time 
to form molecules. 

Equations. — Equations are chemical reactions represented by sym- 
bols and formulae, the substances reacting being placed on 
the left of the equal sign, while the products are placed on 
the right. 

Valence. — Valence is the capacity of atoms of elements to com- 
bine with other atoms. Valence is measured by the number 
of hydrogen atoms one atom of an element can combine with 
or replace. Valence can be changed by oxidation and reduc- 
tion. 

REVIEW QUESTIONS 

1. Upon what facts does the atomic theory rest? 

2. Of what use is this theory? 

3. Give some advantages derived from the use of symbols 
and formulae. 

4. How do formulae differ from symbols? 

5. How would you write the expression, "three molecules of 
sulphuric acid"? 

6. How many hydrogen atoms are there in Ca(OH) 2 ? 

7. Express in words: Ca(OH) 2 + C0 2 = H 2 + CaC0 3 . 

8. What is meant by the nascent state? 



48 A PRACTICAL CHEMISTRY 

9. How does the nascent state explain the oxidizing power 
of hydrogen peroxide 1 ? 

10. What is the difference in the meaning of 2H and Ha? 

11. What relation is there between the atoms on the two sides 
of a chemical equation? 

12. Why can you not make up these equations for yourself? 

13. Show that atoms of all elements do not have the same 
power of combining with other atoms. 

14. How is the valence of an element measured? 

15. What connection is there between valence and the nam- 
ing of compounds? 

16. What new meaning is given in this chapter to the terms 
oxidation and reduction? 

17. Illustrate this. 



CHAPTER VII 
SALT 

Sources of Salt. — No substance is more familiar to the 
ordinary student than common salt. It is found in small 
quantities nearly everywhere, and in large quantities in 
sea water, in the water of salt springs and salt wells, 
and as rock salt in the earth. In various parts of the 
earth salt is obtained from these sources. The rock salt 
is generally dug from the earth and broken and ground 
up fine, but in some places wells are drilled down into 
the salt rock, water is forced down, which dissolves the 
salt, and is then pumped up into tanks. The water is 
then evaporated. This is sometimes done by letting the 
sun shine on the tanks, but more often the brine is heated 
in pans over a fire. Since salt is more soluble than many 
of the impurities which it contains, these are deposited 
before the salt begins to crystallize. So soon as the salt 
crystals begin to form, the brine is drawn off into pans 
and the crystallization continues. The salt is thus ob- 
tained free from most of the impurities. In some places 
natural brines are used in the same way as described 
above, and in others sea water is put through a similar 
process. 

The salt crystals are colorless or white, and of cubical 
shape. When pure, salt does not absorb moisture from 
the air, but usually it contains impurities, such as salts of 
calcium and magnesium, which become damp. This ex- 
plains why salt becomes hard in sifters. 

49 



50 



A PRACTICAL CHEMISTRY 



C2r >>^ 



The quantity of salt required by our country each year 
is very great. Besides its uses as a food and as a pre- 
servative, it is much used in manufacturing, and to some 
extent as a fertilizer. Hydrochloric acid and sodium 
hydroxide are both made from it. It is also our most 
important source of chlorine. 

Composition. — The composition of common salt is of 
interest to us. It contains two important elements, 
sodium and chlorine. Its formula is NaCl. Sodium we 
have mentioned before, and chlorine 
shall soon claim our attention. If a 
solution of salt in water is put into 
a U-shaped tube {see Fig. 19) and 
some indicator, as red litmus, is 
added and an electric current is 
passed through the solution, very 
soon it will be noticed that around 
the negative electrode the litmus 
begins to turn blue, bubbles of gas 
escape from this arm of the tube, 
and if the bubbles are very numer- 
ous they may be lighted and will 
burn as hydrogen. In fact this gas 
is hydrogen. Sodium is being lib- 
erated at the negative electrode, and at once attacks the 
water, liberating hydrogen and forming sodium hydroxide, 
which turns the red litmus blue. In the other tube the 
color of the litmus is soon found to be fading away, and 
bubbles of an ill-smelling, irritating gas are escaping. 
This gas is chlorine. 

The experiment just described is called the electrol- 
ysis of sodium chloride. Any decomposition produced 
by the electric current is called electrolysis. The student 
will recall the electrolysis of water described in the chap- 
ter on water. At Niagara Falls the electrolysis of sodium 




Fig. 19. — Electrolysis of 
Sodium Chloride. A so- 
lution of common salt is 
put into the tube, U, and 
the electrodes, E, E, are 
put into place. 



SALT 



51 



chloride is carried on extensively for the purpose of ob- 
taining chlorine and sodium hydroxide. 

A diagram of the apparatus used is given in Figure 20, 
which the student should carefully study and reproduce 
in his notebook. The apparatus is arranged so that the 
sodium hydroxide is kept separate from the salt solution. 




Fig. 20.— Apparatus for the Manufacture of Sodium Hydroxide by Electrol- 
ysis. Sufficient mercury, H, is put into the bottom of the tank to fill it up a little 
above the lower ends of the partitions, thus separating the tank into three parts 
and preventing the brine in B from mixing with sodium hydroxide solution in W. 
The chlorine is attracted to the positive electrodes, A, A, and escapes through 
C, C. The sodium which is repelled from A, A, dissolves in the mercury and is 
attracted to the cathode, K. This is helped by a rocking motion given to the box 
by means of the eccentric, E, and the fulcrum, F. At K, the sodium reacts with 
the water, making sodium hydroxide and liberating hydrogen which escapes 
through D. 



The chlorine and hydrogen are also carried off by differ- 
ent pipes. The chlorine is used to make bleaching pow- 
der. The sodium hydroxide solution is evaporated to 
dryness, and the sodium hydroxide, which is a white 
solid, is obtained. It is a very strong caustic base, often 
called caustic soda. It is very deliquescent. 

Sodium hydroxide has many uses, the most important 



52 



A PRACTICAL CHEMISTRY 




of which are in connection with the 
refining of oils and in soap making. 

Sodium. — Sodium hydroxide is also 
used as a source of sodium. The hy- 
droxide is dried and melted and an 
electric current is then passed through 
it. ■ The electricity decomposes the 
sodium hydroxide into the three ele- 
ments, sodium, oxygen and hydrogen, 
as shown in Figure 21. Sodium is a 
bright, silver-like metal, not quite so 
heavy as water. It is quickly attacked 
by oxygen and many other elements, 
by water, and by acids. It is not acted 
upon by hydrogen, and may be kept 
bright if sealed up in a tube of this 
gas. Ordinarily sodium is kept in 
naphtha or kerosene, which protect it 
fairly well from the air. When burned 
in the air sodium forms sodium perox- 
ide, a compound now used as a source 
of oxygen. Sodium is also used in 
making sodium cyanide, which is em- 
ployed in extracting gold from its ore, 
in gold and silver plating, and in other 
metallurgical work. 

Sodium reacts with acids in the 
same way that zinc does, replacing 
hydrogen and forming sodium salts, 
but these salts are best made from 
sodium hydroxide and the acids. "We 
have sodium salts of many acids, many 
of which are used in every-day life. 
Chlorine. — We have generated chlorine from salt by 
electrolysis. It may also be obtained by heating a mix- 



Fig. 21. — Appakatus for 
Preparation of So- 
dium by Electroly- 
sis: Castner's Proc- 
ess. The large bottle- 
shaped vessel, B f is 
nearly filled with fused 
sodium hydroxide 
which is kept in the 
fused condition by be- 
ing heated on the out- 
side. The negative 
wire is joined to a 
large carbon rod, C, 
which passes up from 
the bottom of the ap- 
paratus and consti- 
tutes the cathode. The 
anode is a large iron 
cylinder, A. Sodium 
and hydrogen are lib- 
erated at C and escape 
into the can, T, where 
the hydrogen serves to 
prevent oxidation of 
the sodium until it 
has been removed by 
means of a dipper. 
Oxygen is liberated at 
the anode. 



SALT 53 

ture of salt, sulphuric acid, and manganese dioxide, or a 
mixture of hydrochloric acid and manganese dioxide, or 
hydrochloric acid and any strong oxidizing agent. Hy- 
drochloric acid, as the name indicates, is a compound of 
hydrogen and chlorine; strong oxidizing agents oxidize, 
the hydrogen and set the chlorine free. These reactions 
may be expressed in equations thus : 



2HC1 


+ 





= 


H 2 + 


Cl 2 


Hydrochloric 




Oxygen 


from 


Chlorine 


acid 




an oxidizing 










agent 









Or, when salt, sulphuric acid, and manganese dioxide are 
used, 

2NaCl + 2H 2 S0 4 + Mn0 2 = 2H 2 + Na 2 S0 4 

Sodium Manganese Sodium 

chloride dioxide sulphate 

+ MnS0 4 + Cl 2 . 

Manganese 
sulphate 

Chlorine is a poisonous gas with a very unpleasant, 
irritating odor. Its color is yellowish green. It is nearly 
two and one-half times as heavy as air. It should not 
be collected as we do oxygen and hydrogen, since it is 
soluble in water. It may be collected, however, by the 
direct displacement of air. (See Fig. 22.) The delivery 
tube is run down to the bottom of the jar in which the 
chlorine is collected and forces the air out. Chlorine 
is a very active gas, combining violently with many ele- 
ments; for example, if heated metals are thrust into 
chlorine they will burn brilliantly. The compounds 



54 



A PRACTICAL CHEMISTRY 




formed are called chlorides. A jet of hydrogen will also 
burn in chlorine ; the chloride of hydrogen thns formed 

is called hydrochloric acid. 
Chlorine has a great attraction 
for hydrogen and will take it out 
of many compounds. Another 
important property of chlorine is 
that it will bleach many kinds of 
coloring matter. 

Bleaching Powder. — The 
chlorine which is generated in 
large quantities in the electroly- 
sis of sodium chloride is some- 
times condensed into liquid 
chlorine and shipped in steel cyl- 
inders, but more often it is made 
into bleaching" powder by pass- 
ing the chlorine into slaked lime. 
The chlorine unites with the lime 
and forms a complicated com- 
pound which will give off chlorine 
when mixed with acids. This is 
the compound which you have 
bought at the drug store under 
the name of "chloride of lime." 
It is often used as a disinfectant 
and deodorizer, but its chief use 
is in bleaching cotton goods. 

Bleaching Cotton Goods. — I f 
you have ever been in the cotton 
fields and seen the white cotton, 
you would scarcely expect the 
cloth made of cotton to have a yellowish brown color, but 
such is its condition until properly bleached. The bleach- 
ing is quite a long process. The strips of cloth are boiled 



Fig. 22. — Generating and Col- 
lecting Chlorine. In the 
flask, H, a mixture of mangan- 
ese dioxide and hydrochloric 
acid is heated carefully. It is 
important that the mixture be 
well shaken before the heating 
begins. The chlorine, on ac- 
count of its great weight, will 
displace the air in the cylin- 
der, C. The cy Under is nearly 
closed by a glass plate, the 
under side of which is coated 
with vaseline. The jar when 
nearly full of chlorine may be 
replaced by another and at 
the same time sealed with the 
vaselined plate. The flask, H, 
should be protected by a wire 
gauze and should not be 
heated hot enough to cause 
the contents to boil. 



SALT 



55 



in kettles and reeled back and forth through a number of 
tanks and baths containing various substances. The cloth 
is thus thoroughly washed with alkalies, acids, and water, 
after which it is passed through a solution of bleaching 
powder, and then through a dilute acid to decompose the 
bleaching powder and liberate chlorine. The chlorine 




Fig. 23. — Diagram of Bleaching. A is a roll of unbleached goods. B, B are guide 
rollers. P are heated metal plates for singeing off long threads and nap. C and D 
are wash tubs; E and G contain bleaching powder solution; F and H, dilute acid. 
I is the anti-chlor (i. e., a solution to remove chlorine). K is a wash tub, L contains 
starch water. M is a drying roller heated by steam and N, N serve to iron the 
goods smooth. O is a roll of finished goods. 



bleaches the coloring matter. All chlorine must next be 
removed by chemicals and by washing. Should any 
chlorine remain it would soon injure the cloth. A chem- 
ical substance used to remove the chlorine is called an 
"antichlor." The cloth is then starched and ironed. 

In Figure 23 w T e have attempted to show diagrammati- 
cally the various operations of bleaching. 

Hydrochloric Acid. — We have learned by this time 
that salt is a chloride, and, like other chlorides, may be 



56 A PRACTICAL CHEMISTRY 

made by burning sodium in chlorine. But it is also a 
salt, and its name tells us, if we will recall what we have 
learned about the naming of salts, that it is derived 
from hydrochloric acid. This acid contains hydrogen 
and chlorine. Salt contains sodium and chlorine, and 
if we could put hydrogen in place of sodium in the salt 
we should expect to have hydrochloric acid. The reac- 
tion is expressed by this equation, if the process is con- 
ducted at a high temperature : 

2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1, 
Sodium 
sulphate 

but when the mixture is not heated the reaction is : 

NaCl + H 2 S0 4 = NaHSO, + HC1 
Acid sodium 
sulphate. 

The hydrochloric acid comes off as a heavy, irritating, 
colorless gas having all the properties of acids. Hydro- 
chloric acid gas is very soluble in water and will even 
condense the moisture of the air. A water solution 
containing about 40 per cent of acid is the hydrochloric 
acid of commerce. It is sometimes called muriatic acid. 
In the factory the mixture of salt and sulphuric acid is 
heated in iron vessels so arranged that the hydrochloric 
acid gas passes into large absorption towers filled with 
coke and other acid-resisting material over which water 
trickles. Here the gas is dissolved in the water and falls 
to the bottom, where it is collected and bottled. 

A number of other acids may be made by a process 
similar to this, that is, by heating their salts with sul- 
phuric acid and driving off the required acid as a gas or 
vapor. 



SALT 57 



SUMMARY 



Salt. — Salt occurs in natural brines and as rock salt in the earth. 
It is obtained by evaporating these brines and from salt 
mines. It is composed of sodium and chlorine. 

Chlorine and Sodium Hyaroxide. — Chlorine and sodium hydrox- 
ide are obtained from salt by electrolysis. Chlorine is a 
heavy, greenish yellow, gaseous element with an irritating 
odor. It is much used in bleaching. Sodium hydroxide is 
a strong base called caustic soda, much used in refining oils 
and making soap. Sodium is obtained by the electrolysis of 
sodium hydroxide. It is a silver-like metal which is acted 
upon by air, moisture, carbon dioxide, and all acids. It is 
kept in some inactive gas or liquid. Chlorine may be gen- 
erated by oxidizing hydrochloric acid and collected by the 
displacement of air when wanted for laboratory purposes. 

Hydrochloric Acid. — Hydrochloric acid is a heavy, irritating, 
colorless, acid gas which dissolves freely in water, forming 
a solution which may contain as much as 40 per cent of 
hydrochloric acid. 

REVIEW QUESTIONS 

1. Describe the process of getting salt from each of its 
sources. 

2. Why does salt often become moist when exposed to the air? 

3. What are some of the uses for which salt is required? 

4. How would you show that salt contains chlorine? 

5. Why was litmus added to the salt solution in the electrol- 
ysis of sodium chloride experiment? 

6. State the properties of sodium hydroxide. 

7. Name some of its uses. 

8. Why can you not get sodium by the electrolysis of a 
solution of sodium hydroxide? 

9. What is sodium cyanide? 

10. How is oxygen generated from sodium peroxide? 

11. Why cannot chlorine be collected over water? 



58 A PRACTICAL CHEMISTRY 

12. Describe some striking experiments with chlorine. 

13. How is bleaching powder made? 

14. Why does "chloride of lime" give off chlorine when ex- 
posed to the air. 

15. How can you generate chlorine from bleaching powder? 

16. How is hydrochloric acid made from salt? 

17. Is a solution of hydrochloric acid heavier than water? 

18. Why does hydrochloric acid "fume" in moist air? 

19. How would you make nitric acid? 

20. Write an equation to represent a method of generating 
chlorine. 



CHAPTER VIII 
THE AIR 

The great body of mixed gases surrounding the earth 
is often called the atmosphere. Any portion of this at- 
mosphere is called air. Much of our comfort depends 
upon the condition of the air which surrounds us ; hence 
the importance of a study of the atmosphere. 

Weight of the Atmosphere. — Although man has never 
ascended far above sea level, he has found through scien- 
tific calculations that the air must extend many miles 
above the earth. The highest parts of the atmosphere 
must be quite rare, since the air is not so dense and 
does not press down with nearly so much force upon the 
tops of high mountains as it does at sea level. By care- 
ful measurements it is shown that this air pressure varies 
from day to day, the average pressure at sea level being 
about 15 pounds per square inch. The pressure of the 
air is due to its weight. Air pressure is measured by 
the barometer, the simplest form of which is a glass tube, 
closed at one end, filled with mercury and inverted in 
a cup of mercury. The height of the mercury in this 
tube will vary with the air pressure, but it averages about 
760 millimeters, or nearly 30 inches in height. We do 
not feel the pressure of the air because it is on all sides 
and our bodies are adjusted to sustain it. Since, as we 
have said, the pressure becomes less as we go higher up 
into the atmosphere, the mercury in a barometer gradu- 
ally sinks lower and lower while the barometer is being 

59 



60 



A PRACTICAL CHEMISTRY 



carried upward. In this way it is possible to measure 
the height of mountains. 

Moisture in the Air.— It is sometimes noticed that 
moisture collects on the outside of an ice water pitcher 
when standing in a warm room. The quantity of this 

moisture is much greater at 
some times than at others. 
This means that the quantity 
of moisture in the air is not 
always the same. Generally 
there is not so much moisture 
in the air as it can hold, and 
the quantity which it can hold 
varies with the temperature, 
increasing as the temperature 
increases and becoming less as 
the temperature decreases. 
The quantity of moisture 
which the air contains as com- 
pared with what it can hold at 
that temperature is called the 
relative humidity. Relative 
humidity is usually expressed 
in per cent ; thus, if the air at 
a certain temperature contains 
one-half as much moisture as 
it is capable of holding at that 
temperature, the relative hu- 
midity is said to be 50 per cent. If the air is cooled, the 
relative humidity is increased until saturation is reached, 
and after that, if the cooling is continued, some of the 
moisture must drop out as dew, fog, rain, or snow. Thus 
we see that at least three things which affect the condition 
of the air are constantly changing — its temperature, pres- 
sure, and relative humidity. These three factors have 




Fig. 24. — The Principle op the 
Barometer. The tube, B, is 
filled with mercury and inverted in 
the cup of mercury. The mercury 
in the cup is forced by the pres- 
sure of the air up into the tube to 
such a height that the weight of 
the mercury column balances the 
weight of an air column of equal 
cross-section. A similar apparatus 
supplied with a suitable frame and 
graduated scale constitutes a ba- 
rometer. 



THE AIR 



61 



p m 

U3^ 



much to do with making up the climate of a country. The 
air receives its moisture by the evaporation of oceans, 
rivers and other water, as we have previously mentioned. 
Its heat comes largely from contact with the earth, which 
is heated by the sun's rays. The differ- 
ences in pressure are due to many 
causes. 

Composition. — We have called air "a 
mixture," and we have mentioned that 
it contains oxygen. Mixed with the 
oxygen are many other gases, some of 
which we shall study. In 100 volumes 
of dry air there are about 21 volumes 
of oxygen and 78 volumes of nitrogen 
and one volume of rare elements. 

Nitrogen. — If we were to float some 
burning phosphorus on a cork on water 
and turn over it a large jar of air, the 
phosphorus would continue burning for 
some time until practically all of the 
oxygen in the jar had been used up. 
The jar would soon be filled with a 
white oxide of phosphorus. This oxide 
would dissolve in the water and leave a 
considerable volume of gas in the jar. 
The remaining gas is nearly pure nitro- 
gen. Nitrogen may also be obtained by 
decomposing certain of its compounds. 
From your own experience with air 
nitrogen has neither color, taste, nor 



Fig. 25. — The Volume 
of Oxygen in Air. 
A definite volume of 
air (as 100 c.c.) is 
measured into the 
cylinder, C, which is 
inverted in a jar of 
water. The bit of 
phosphorus, P, is 
thrust up into the 
tube by means of a 
wire and the whole 
allowed to stand for 
24 hours. The vol- 
ume of remaining 
gas is then read. 
The difference be- 
tween this reading 
and original volume 
of air represents the 
volume of oxygen 
removed. 



you know that 
smell. Experi- 
ments have shown it to be a little lighter than air, since 
it is 14 times as heavy as hydrogen, while air is 14.4 
times as heavy. The character of nitrogen is almost the 
opposite of that of oxygen. Oxygen, you will recall, 
combines with almost everything, and in combining often 



62 



A PRACTICAL CHEMISTRY 




gives out much energy as heat and light. Nitrogen com- 
bines directly with but few elements, and often requires 
much energy to cause its combination. Nitrogen in com- 
bination, however, may be easily changed into other com- 
pounds, hence the number of nitrogen compounds is very 
great, and many of them are of much importance. Some 
of them are necessary for the growth of plants and 
animals. 

In the air nitrogen dilutes the oxygen 
and retards its action. This is very de- 
sirable, for we have seen how violently 
many things burn in pure oxygen. If 
the air were pure oxygen, our stoves and 
furnaces would be consumed and fire 
could be controlled only with the great- 
est difficulty. 

Products of Combustion in the Air. 
— We have had occasion under oxygen 
to learn something of combustion, and 
have learned to think of ordinary com- 
bustion as oxidation or the formation of 
oxides accompanied with the evolution 
of light and heat. The oxides formed in 
ordinary combustion are mostly gases and escape into the 
air. If we lower a candle into a flask of air (see Fig. 
26) we notice that the candle flame becomes less and less 
bright and soon goes out entirely. This is as we should 
expect it to be, knowing that the oxygen in the flask is 
used up as the candle burns. But other questions sug- 
gest themselves. "What oxides are formed by the burn- 
ing of the candle, and how may we detect them? Let us 
continue our experiment by first removing the candle 
from the flask and then throwing in a little lime water 
and shaking the contents of the flask. At once the lime 
water becomes milky. If we repeat the experiment by 



Fig. 26. — Combustion 
Produces Carbon 
Dioxide. The can- 
dle, C, is lowered by 
the wire, W, into the 
flask. It is removed 
in the same way be- 
fore adding the lime 
water. 



THE AIR 63 

burning carbon in pure oxygen, thus making carbon 
dioxide, the lime water again becomes milky. In other 
words, our test for carbon dioxide is the formation of 
this milky appearance in lime water when shaken up 
with a gas. The milky appearance is due to the forma- 
tion of a precipitate of calcium carbonate. Equations 
representing these reactions are : 

(1) C + 2 := C0 2 , 

Carbon Carbon 
and dioxide 
oxygen 

which represents the making of the carbon dioxide : 

(2) C0 2 + H 2 = H 2 C0 3 , 

Carbonic 
acid 

which shows the formation of an acid from an oxide dis- 
solving in water; 

(3) H 2 CO ;J + Ca(OH) 2 = 2H 2 + CaC0 3 , 

Calcium Calcium 

hydroxide carbonate 

which represents the neutralization of the acid and base 
and the formation of the white precipitate of calcium 
carbonate. By holding a flask of cold water above a 
burner flame we at once notice moisture collecting on 
the flask. This shows that water (hydrogen monoxide) 
is also a product of ordinary combustion. These two 
substances, moisture and carbon dioxide, are the most 
important products of combustion which escape into the 
air, but smaller quantities of other gases, particularly 
sulphur dioxide and carbon monoxide, are often given off 
from fires. 



64 A PRACTICAL CHEMISTRY 

Carbon Dioxide a Product of Respiration. — By the aid 

of the lime water test it may easily be shown that carbon 
dioxide is given off in the breath. The continuous breath- 
ing of many people in a room soon increases the amount 
of carbon dioxide in the air of the room. But carbon 
dioxide, although it will suffocate people if taken in 
large quantities, is not a poison. Why, then, is ventila- 
tion so necessary? The air we breathe out from our 
lungs carries with it besides the carbon dioxide quantities 
of decomposing tissue and bacteria, which soon produce 
the unpleasant odor which one notices on entering a 
poorly ventilated room. By measuring the quantity of 
carbon dioxide in the room we may estimate the quantity 
of other impurities in the air, since these impurities will 
increase at about the same rate as the carbon dioxide 
while the air of the room is being breathed over and 
over. It is easy to imagine how injurious to health 
such impurities are, and too much emphasis cannot be put 
upon the importance of pure air. Carbon dioxide enters 
the air in small quantities from many other sources. 
Among these we may mention gases escaping from vol- 
canoes and cracks in the earth, from mineral springs, 
from fermentation and decay, and other chemical changes. 
Carbon Dioxide Removed from the Air. — We might 
well expect from what has been said above that carbon 
dioxide is all the time accumulating in the air, but this 
is not the case ; in much the same way that animals- 
breathe in oxygen and give out carbon dioxide, so do 
plants in the daytime take in carbon dioxide and give 
out oxygen. Thus the supply of oxygen is renewed and 
the proportion of carbon dioxide is kept down to about 
4 parts in 10,000 parts of air. As we continue our study 
of chemistry we shall find a number of other ways in 
which carbon dioxide in small quantities is removed, from, 
the air. 



THE AIR 



65 



Rare Elements in the Air. — In addition to the sub- 
stances thus far mentioned, the air has been found to 
contain small quantities of a number of rare gases. These 
gases are elements which show no chem- 
ical properties, and form, so far as we 
know, no compounds. Their physical 
properties have been studied and the 
various gases have received names as 
follows: argon, helium, neon, krypton, 
and xenon. 

Everyone has noticed that the air 
usually contains more or less dust. 
These dust particles serve several pur- 
poses, all of which are of some impor- 
tance. It is evident to all that bacteria 
are often carried on dust particles. Part 
of the diffused light which enters our 
rooms is reflected in by particles of dust. 
In the formation of raindrops it is neces- 
sary that the drop have something to 
begin to form around; dust particles 
may often serve for this purpose. The 
same thing is true in the formation of 
snow crystals. 

Liquid Air. — Although some gases 
were many years ago converted into 
liquids, the air and hydrogen were 
looked upon as "permanent" gases be- 
cause they could not be liquefied by 
pressure at the ordinary temperature. 
In more recent years, however, the lique- 
faction of air has become a simple mat- 
ter, owing to the discovery that it must be made very cold 
as well as compressed. The most satisfactory way to ac- 
complish this is to compress the air by means of a pump 




Fig. 27. — Action of 
Gkeen Plants up- 
on Carbon Diox- 
ide. Through the 
action of light on 
the green leaves in 
the flask, oxygen is 
liberated which may 
be collected in the 
inverted test tube 
in the funnel above. 
This illustrates how 
plants return oxy- 
gen to the air in 
place of the carbon 
dioxide which is fur- 
nished to them by 
the air. 



66 A PRACTICAL CHEMISTRY 

until a very great pressure is reached. This process heats 
the air greatly, and its temperature is again reduced to 
that of ordinary air by allowing cold water to flow over 
the pipes in which the air is being compressed. The com- 
pressed air next passes into a coil of pipe which is sur- 
rounded by a cylinder. At the end of this coil of pipe is a 
hole the size of a needle, through which the compressed air 
escapes into the cylinder, which in turn is open to the out- 
side air. When a compressed gas is thus suddenly allowed 
to expand it absorbs much heat. The compressed air com- 
ing through the needle hole mentioned above takes heat so 
rapidly from the coil of pipe within the .cylinder that 
this pipe is soon reduced to the temperature at which air 
becomes a liquid, and liquid air begins to drip from the 
coil into the cylinder below. The temperature of liquid 
air is nearly 200° C. colder than that of melting ice. Ow- 
ing to the fact that the nitrogen of liquid air evaporates 
and becomes gaseous nitrogen much faster than the oxy- 
gen evaporates, liquid air is always very rich in oxygen ; 
in fact, after standing a while it becomes as much as 
90 per cent oxygen. On this account combustion in 
liquid air is extremely violent. The low temperature of 
liquid air makes it useful for freezing other things, and 
for the cooling of other gases. Mercury and alcohol may 
be easily frozen with liquid air. 



SUMMARY 

The Atmosphere. — The atmosphere is a mixture of gases which 
extends for many miles beyond the earth. Owing to its 
weight, it exerts a pressure of about 15 pounds per square 
inch upon the surface of the sea. This pressure is repre- 
sented by a column of mercury 760 millimeters high. 

Relative Humidity. — The relative humidity is the quantity of 
moisture which the air contains as compared with what it 



THE AIR 67 

can hold at that temperature. There are approximately 78 
volumes of nitrogen, 21 volumes of oxygen, and one volume 
of rare elements in 100 volumes of dry air. 

Nitrogen is a colorless, odorless, tasteless gas which enters 
into combination with difficulty, but when once combined 
with other elements passes easily from one compound to 
another. 

Carbon dioxide is a product of combustion and respira- 
tion and is present to the extent of 4 parts in 10,000 parts of 
air. It may be detected by the lime water test. Small quan- 
tities of inactive rare elements and other gases are found 
in the air. 
Liquid Air. — Liquid air is made by compressing and cooling air. 
On account of the rapid evaporation of the nitrogen, liquid 
air contains a large percentage of oxygen. Its temperature 
is about 200° below the Centigrade zero. 



REVIEW QUESTIONS 

1. Give three reasons for thinking that air is a mixture. 

2. What uses of air have been nearly discarded? 

3. What is the atmosphere? 

4. How is its weight measured? 

5. Under what conditions can you feel the weight of the air? 

6. How can we show that the air contains moisture? 

7. Define relative humidity. 

8. How does the relative humidity affect breathing? 

9. . State the composition of dry air. 

10. How may we obtain nitrogen from the air? 

11. In what other way can we get nitrogen? 

12. Compare nitrogen with oxygen. 

13. What results would follow if the air became pure oxygen ? 

14. How can you detect the products of combustion in the air? 

15. Why are you not injured by the carbon dioxide from 
soda water while drinking it? 

16. Why does the air of an unventilated room become bad? 

17. Does carbon dioxide accumulate in the atmosphere? 



68 A PRACTICAL CHEMISTRY 

18. Recount the sources of the carbon dioxide of the air. 

19. Would you find more carbon dioxide in the air of the 
fields or of the city? 

20. Follow the cycle of carbon in nature. 

21. How do the rare elements in the air differ from the other 
elements thus far studied? 

22. What is known of these elements'? 

23. Give their names. 

24. What is the importance of dust in the air? 

25. How may air be liquefied? 



CHAPTER IX 
CARBON AND FUELS 

In a previous chapter we have spoken of combustion 
and of some of the products of combustion. Let us now 
give some thought to fuels, the substances burned. We 
may include under this head wood, charcoal, coal, coke, 
petroleum, acetylene, water gas, coal gas, etc. We wish 
to examine each of these, learn its character, composi- 
tion, and properties. But first of all we come upon the 
fact that they all contain carbon, and this brings us to 
a study of that element as preliminary to a study of fuels. 

Forms of Carbon. — Like oxygen, carbon exists in more 
than one form. Each has its own peculiar structure and 
properties, but all on burning yield carbon dioxide ; all 
must therefore contain carbon. These forms are known 
as diamond, graphite, and amorphous carbon. 

Diamond. — In some parts of the earth's crust, crystals 
of carbon are found. These are called diamonds. They 
vary greatly in size and color. The finest specimens are 
colorless, others are yellow, green, blue, and black. The 
colorless diamonds, and also some of the colored varieties, 
when properly cut and polished, are very brilliant and 
are highly valued as gems. The diamond is the hardest 
of all minerals, and on this account has been found valu- 
able as a point for drills for drilling through rocks. 
Black diamonds are used for this purpose. Black dia- 
monds are also set in the teeth of saws for sawing stone. 
They are also being used as needle points in phonographs. 



70 



A PRACTICAL CHEMISTRY 



Of late years a large part of the world's diamond supply 
has , come from South Africa. The diamond is heavier 
than other forms of carbon; its specific gravity is 3.5, 
while that of graphite is 2.15. The specific gravity of 
amorphous carbon is variable. Very small diamonds 
have been made by dissolving carbon in molten iron in 
an electric furnace, and then suddenly cooling the iron 
by plunging it into cold water. The iron on the outside 
is thus quickly hardened, and by its contraction produces 
a great pressure upon the carbon within. It seems prob- 




Fig. 28. — Graphite Crucibles. Graphite is well suited for this purpose because of 
its high melting point and high kindling temperature. 



able, therefore, that diamonds were formed in the earth 
from melted carbon under great pressure. 

Graphite. — The graphitic form of carbon is very differ- 
ent from the other two forms. In color it is black with 
a gray luster. It has a sort of crystalline structure, and 
is made up of thin flakes which overlap one another and 
give it something of the appearance of ancient writing. 
The particles of graphite are very hard, but they slide 
over one another so easily that graphite is often used to 
lubricate machinery. On account of the high tempera- 
ture which graphite will withstand, it is much used for 
making crucibles in which to melt metals. For the 



CARBON AND FUELS 71 

"leads" of lead pencils graphite is mixed with clay and 
pressed into rods. These rods are placed in grooves in 
a thin board and a similarly grooved board glued on top 
of them, after which the pencils are sawed apart and 
rounded into shape. The hardness of the pencils is regu- 
lated by the percentage of clay mixed with the graphite. 
Graphite is found useful in the manufacture of many 
other things, among which may be mentioned stove pol- 
ish, paint, and electrodes. Although graphite is found 
and mined in many parts of the world, large quantities 
of it are now made artificially. In this process carbon 
is heated with a small percentage of iron in the electric 
furnace. 

Amorphous Carbon. — Under this heading we must in- 
clude all free or uncombined carbon except the diamond 
and graphite. As the name indicates, amorphous carbon 
does not have crystalline form. It is black in color and 
is known by various names, as coke, coal, charcoal, etc., 
which we shall take up in detail. 

Wood. — Of all the various forms of fuel used by man 
at the present time none has been longer known and used 
than wood. It might for this reason seem proper to be- 
gin our study of fuels with a study of wood, but there 
is a deeper reason why it should come first — nearly all. 
fuels are derived directly or indirectly from wood. When 
obtained from different kinds of trees, wood varies some- 
what in appearance and composition, but always contains 
cellulose, water, and mineral salts. Cellulose is a definite 
chemical compound of carbon, hydrogen, and oxygen. 
It is white in color, insoluble in most liquids, but reacts 
with several chemicals. Some of its compounds we shall 
have occasion to mention later. Cotton and paper are 
nearly pure cellulose. 

During combustion the carbon, hydrogen, and oxygen 
of wood are converted into carbon dioxide and water. 



72 A PRACTICAL CHEMISTRY 

The mineral matter remains as ash, commonly known as 
wood ashes. 

Charcoal. — If wood is heated with, little or no air, its 
hydrogen, oxygen, and some of its carbon are converted 
into volatile compounds which pass off as gas or vapor, 
leaving the mineral matter and the rest of the carbon 
behind as a porous black mass known as charcoal. Such 
a process as this is called destructive distillation. The 
volatile substances when condensed are found to include 
wood alcohol, acetic acid, acetone, and tar. "When it is 
desired to save the volatile products the wood is heated 
in retorts and the products condensed and collected. 
Charcoal is often made in kilns, and the volatile sub- 
stances permitted to escape into the air. 'A kiln may be 
defined as an oven in which substances are heated for 
the purpose of removing the volatile matter which they 
contain. A retort is a closed vessel for the same purpose 
as a kiln, but supplied with an exit pipe through which 
the volatile substances may be conducted to suitable con- 
densers. Sometimes the wood is simply piled on the 
ground. It is then well covered over with brush, sand, 
and sod to exclude the air. A fire is then started and 
enough air admitted to allow the wood to smolder for 
some time, thus driving off the volatile substances and 
converting the wood into charcoal. When the process is 
complete the fire is extinguished and the kilns allowed 
to cool. In place of the temporary kilns of earth just 
described, dome-shaped ones of a more substantial char- 
acter are often used. These are provided with a door 
at one side to receive the wood and an exit pipe near the 
top for the volatile products. At the bottom there is a 
grate and an ashpit with doors that may be closed air- 
tight. The production of charcoal with a kiln of this sort 
is very similar to its manufacture with an earth kiln. 

Charcoal is a porous black substance with great capac- 



CARBON AND FUELS 73 

ity for absorbing gases, coloring matter, etc. On this 
account it is much used for filters. As a fuel it furnishes 
much heat with little or no flame and smoke. The prod- 
ucts of its combustion are carbon dioxide, carbon mon- 
oxide, and ash. Charcoal is also used as a reducing 
agent in metallurgy. 

Boneblack — Animal Charcoal. — Boneblack and animal 
charcoal are forms of charcoal made by the destructive 
distillation of bones and of blood mixed with sodium 
carbonate. They are used for filtering sugar solutions in 
the refining of sugar. 

Coal. — The surface of the earth has passed through 
many changes during the past geological ages. Climate 
also has been very different at different times. There 
have been times when large areas which are now moun- 
tains were swamps in which grew in abundance many 
plants similar to those now growing in tropical regions. 
The leaves, branches, and trunks of these plants fell into 
the water of the swamps and underwent a slow decom- 
position like that which takes place in peat bogs. This 
continued through long periods, the swamps at last be- 
came filled up, and through various geological changes 
the surface became elevated, while the vegetable matter 
thus stored away slowly hardened into coal. Evidence 
that coal was at one time vegetable matter is found in 
the fossils which abound in the coal beds. It is also 
often evident from the structure of the coal as shown 
under the microscope. The many kinds of coal known 
may be so arranged as to show all the various stages 
which coal has passed through during its formation. 

The various kinds of coal are usually spoken of as 
hard, or anthracite coal, and soft, or bituminous coal. 
There are many varieties of each. In general it may be 
said that anthracite contains a large percentage of car- 
bon and but little hydrogen, while soft coals contain 



74 A PRACTICAL CHEMISTRY 

much hydrogen. Coal usually contains some sulphur, 
nitrogen, oxygen, and a considerable amount of mineral 
matter, which forms the ash when the coal is burned. 

Combustion of Coal. — In the ordinary combustion of 
hard coal in furnaces with plenty of air the carbon is 
converted into carbon dioxide, the hydrogen into water, 
and the sulphur becomes sulphur dioxide and hydrogen 
sulphide. The nitrogen escapes as free nitrogen. When, 
however, the air supply is not sufficient much carbon 
monoxide is formed. This is a colorless, odorless, taste- 
less, poisonous gas which is often seen burning with a 
blue flame over the surface of coal fires. The coal in the 
lower part of the fire burns to carbon dioxide, which as 
it passes up through the hot coal above is reduced to 
carbon monoxide. If sufficient air is present, the carbon 
monoxide burns to carbon dioxide at the surface of the 
fire, as described. If, on the other hand, the air supply 
is limited, the carbon monoxide escapes without burning, 
and, since the carbon which it contains is but half burned, 
much fuel is thus lost. The fire gases which often escape 
from leaking stoves and furnaces contain much carbon 
monoxide, as well as carbon dioxide, sulphur dioxide, and 
hydrogen sulphide. All of these gases, with the excep- 
tion of carbon dioxide, are poisonous, and their presence 
in the air of dwellings is very objectionable. 

The burning of soft coal is not very different from that 
of hard, except that large quantities of small particles 
of carbon pass up the flue from the soft coal as a heavy 
black smoke. 

Coke. — Just as the destructive distillation of wood 
gives us charcoal, so the same process when applied to 
bituminous coal produces coke. The coal is heated in 
closed retorts or in ovens. During this process the vola- 
tile materials, including ammonia, tar, and coal gas, are 
driven out. All of these are useful and valuable. With 



CARBON AND FUELS 



75 



the old beehive ovens, which have been and still are much 
used in this country, these volatile products are lost. 
These ovens are dome-shaped, with a door near the top. 
A considerable quantity of coal is put into each oven and 
lighted. By thus burning part of the coal the volatile 
matter is driven off from the remainder and either burns 
or escapes into the air. The coke is quenched with water 
and removed from the oven. 
Retort ovens which are so made 
that the gas is collected and used 
to heat the retorts from the out- 
side are now coming into use. 
Thus a considerable amount of 
fuel and ammonia may be saved. 
Coke is a very hard, porous, 
gray-black solid, with metallic 
luster. On account of its poros- 
ity it makes a very hot fire, and 
its hardness enables it to stand 
the great weight of ore in the 
blast furnaces, where it is used 
in large quantities to reduce 
iron. 

Hydrocarbons. — T he gas 
formed in the destructive distil- 
lation of coal is composed largely 

of compounds of carbon and hydrogen. These compounds 
are called hydrocarbons. The hydrocarbons are a very 
large and important class and include some of our most 
common substances. The simplest member of this class is 
marsh gas, or methane, CH 4 . It gets its name from being 
found in marshes, where it is formed by the decomposition 
of vegetable matter under water. When a stick is pushed 
down into the mud at the bottom of a swamp bubbles 
of methane rise and may be collected in an inverted jar 




Fig. 29. — Collecting Marsh 
gas From the Bottom of a 
Pond. The mud at the bot- 
tom of most ponds is mixed 
with decaying leaves; marsh 
gas is formed during this de- 
cay and when a stick is 
pushed into the mud, bubbles 
of the gas escape and may be 
collected as shown in the 
figure. 



76 



A PRACTICAL CHEMISTRY 




of water. Boys sometimes find large bubbles under the 
ice on ponds in winter ; these contain marsh gas. Marsh 
gas burns with a hot, pale flame, and if mixed with air 
explodes violently when lighted. The fire-damp explo- 
sions which occur so often in the coal mines are explo- 
sions of mixtures of marsh gas and air. 
Fire damp is marsh gas, or methane, 
which has collected in a coal mine. 
Methane is an important constituent of 
natural gas, and also occurs in ordinary 
illuminating gases. Natural gas is 
found in many parts of the earth and 
is obtained by drilling wells. The gas 
thus obtained is excellent fuel and is 
piped for considerable distances to fac- 
tories and homes. 

A large number of other hydrocar- 
bons are related chemically to methane, 
and in fact may be indirectly made 
from it. These constitute the methane, 
or marsh, gas series. Other series of 
hydrocarbons are named after some 
prominent member which they contain, 
and from which they may be indirectly 
made. Thus we have the acetylene 
series, the ethylene series, the benzene 
series, and several others. 

Acetylene. — Acetylene, C 2 H 2 , has 
been known for many years, but only 
since the discovery that it can be made from calcium car- 
bide and water, and the perfecting of electric furnaces in 
which the carbide can be made in large quantities, has it 
come into prominence as a substitute for other illuminating 
gases. Figure 100 is a picture of an electric furnace used 
in the manufacture of carborundum. It is similar in con- 



Fig. 30. — Apparatus 
for Generating 
Acetylene. Cal- 
cium carbide is put 
into the cylinder, C, 
and is prevented 
from falling out by 
the wire gauze at 
G. Water enters 
through this gauze 
and reacts with the 
carbide. When the 
gas pressure becomes 
sufficiently strong 
the water is forced 
out again. 



CARBON AND FUELS 



77 



Air 




struction to those used for the manufacture of calcium car- 
bide. In these electric furnaces coal and quicklime are 
heated to a very high temperature by means of the electric 
current. The calcium carbide thus formed is a dark, stone- 
like solid which reacts violently with water, forming acety- 
lene gas and slaked lime (CaC 2 -f- 
2H 2 = C 2 H 2 -[- Ca(OH) 2 . The gas 
has an unpleasant odor, burns with a 
smoky flame when lighted in an or- 
dinary jet, but produces a very 
brilliant light when burned in a spe- 
cial burner constructed on the plan 
of a Bunsen burner, so that the gas 
may be mixed with the proper 
amount of air before coming into the 
flame. Acetylene forms explosive 
compounds with copper, and should 
not be passed through copper or 
brass pipes or fixtures. It also forms 
an explosive mixture with air. 

Acetylene generators are often so 
arranged that the gas is made only as needed. This may 
be accomplished by letting, from time to time, small quan- 
tities of the carbide fall into the water. 

The acetylene torch is an apparatus so constructed 
that a strong blast of oxygen may be forced into an 
acetylene flame. It is used wherever a great heat is to 
be applied to a small surface, as in the welding or fusing 
together of aluminum castings. The acetylene torch is 
also used in the cutting of steel beams, in the repairing 
or wrecking of steel structures. The steel is rendered 
white hot by the acetylene flame, and burned through by 
the large excess of oxygen in the flame. 

The hydrocarbons of the benzene series constitute a 
very important group, but, since they do not enter into 



Fig. 31. — The Acetylene 
Burner. The structure 
of the burner is that of 
two very small Bunsen 
burners set at an angle 
so that their flames 
strike against each other. 
The gas escapes at the 
central holes and mixes 
with the air which en- 
ters as shown in the 
figure. 




Fig. 32. — The Acetylene Torch in Use. When it becomes necessary to remove a 
steel structure, the various parts are cut apart by the acetylene flame. The picture 
represents the removal of railroad tracks over the Harlem River. 



CARBON AND FUELS 79 

fuels so largely as some others, they need not now be 
considered. 

Petroleum. — In many parts of the United States and 
in some other countries large quantities of rock oil or 
petroleum are obtained by drilling holes into the earth. 
The oil often flows freely from these wells for a time, 
but later must be pumped out. It is collected in large 
tanks, measured, and run through long pipe lines to the 
refineries, often located many miles away. The American 
petroleum is composed almost entirely of hydrocarbons 
of the methane series in various proportions. These hy- 
drocarbons are of many different kinds ; some have very 
low boiling points, while the boiling points of others are 
quite high. Some are light liquids, others thick, heavy 
liquids, and still others are solids dissolved in the liquids. 

At the refinery the primary object is to separate the 
petroleum by fractional condensation into a number of 
substances boiling at different temperatures. For this 
purpose the petroleum is put into a large boiler and 
gradually heated. As the temperature of the still rises 
the hydrocarbons are converted into vapors which pass 
over into the condensers. These are so arranged that 
those hydrocarbons with high boiling points, which will 
condense first, may run off into separate receiving tanks. 
The petroleum is thus separated into three or four oils, 
while a considerable amount of coke remains in the 
boilers. These oils are in turn redistilled and separated 
into several oils and solids such as gasoline, benzin, 
kerosene, paraffin oil, vaseline, and paraffin wax. These 
are all further purified before being placed on the mar- 
ket. The purification of the kerosene, which is the oil 
used for illuminating purposes, includes washing it suc- 
cessively with sulphuric acid, water, and sodium hydrox- 
ide. Many thousands of tons of sulphuric acid are used 
for this purpose. The kerosene through this treatment 



80 A PRACTICAL CHEMISTRY 

is not only freed from those impurities which would clog 
the lamp wicks and prevent its combustion, hut the more 
volatile compounds, which would easily catch fire and 
explode in the lamp, are also removed. The law in most 
states requires the kerosene to be subjected to a fire or 
flashing point test. This is done by slowly heating a 
small quantity of the oil and noting the temperature at 
which its vapor will flash when a small flame is brought 
to a certain distance from the oil. 

The fuel uses of such products as gasoline and kerosene 
are very numerous and familiar to all. Much crude 
petroleum is also being used for fuel in' furnaces for 
manufacturing plants and on steamships, s 

Coal Gas. — When the destructive distillation of soft 
coal is conducted in closed retorts, as we have previously 
mentioned while discussing coke, the volatile products, 
including ammonia, coal gas and coal tar, may all be 
saved. This process is carried on at the gas works. The 
apparatus used consists of many parts and may be de- 
scribed as follows : The coal is heated in large retorts 
made of fire clay and arranged in groups, so that a num- 
ber can be heated from one coke fire. About 300 pounds 
of gas coal is put into each retort and the retorts closed 
up so tight that no air may enter. The gases and volatile 
matter escape from the retorts through upright pipes 
which bend over and dip down into a large pipe known 
as the hydraulic main. This main is partly filled with 
water, through which the gas bubbles, leaving much of 
the tar and ammonia behind. The ammoniacal liquor 
and tar pass out of the main into a tar well, while the 
gases are forced in the other direction by a pump into 
the tar extractor, where the tar is removed by friction. 
The tar extractor consists of a number of vertical cylin- 
ders, one within the other. The lower ends of these cylin- 
ders are closed with water, but the sides and upper ends 



CARBON AND FUELS 81 

are perforated with numerous holes. The gas passing 
through these holes loses its tar by friction. It is then 
cooled in the condensers both by air and by water ; the 
gas may pass through pipes surrounded with air, or 
water may pass in pipes in one direction through cham- 
bers in which gas is passing in the opposite direction. 
From the condensers the gas passes to the washers and 
scrubbers for the removal of remaining ammonia. Wash- 
ers are contrivances to make the gas pass in bubbles 
through water. Scrubbers are either perpendicular cylin- 
ders filled with wood over which water is flowing, or 
horizontal cylinders containing large wooden wheels. 
These are kept constantly wet by dipping, as they turn, 
into water in the lower part of the cylinders. In either 
case the gas passing through the scrubbers is brought 
into contact with water, which absorbs the ammonia. 
The gas next goes into the purifiers, which are very large 
iron boxes with shelves carrying iron oxide mixed with 
shavings. Here the sulphur compounds are removed by 
reacting with the iron oxide. After leaving the purifiers 
the gas is measured and passed into large holders, such 
as may be seen in any city. From the holder it passes 
through the mains to the consumers. When purified, coal 
gas contains hydrocarbons, carbon monoxide, and hy- 
drogen. 

Water Gas. — Another kind of illuminating gas much 
used in this country is usually called water gas. The un- 
derlying principle involved in making this gas is that 
highly heated carbon will decompose steam, forming hy- 
drogen and carbon monoxide, according to the equation, 
H 2 -f- C = H 2 + CO. In the factories where water gas 
is made this reaction is accomplished by forcing steam 
through a very hot coke fire. The mixture of hydrogen 
and carbon monoxide thus formed produces much heat, 
but little light when burned. Hence, if it is to be used 



C .2 




CARBON AND FUELS 83 

for illuminating purposes, it must be enriched with hydro- 
carbons obtained from petroleum. When used merely as 
fuel, as in the making of steel, the hydrocarbons are not 
added. 

The manufacture of water gas may be briefly described 
thus: Steam is forced through a very hot coke fire, the 
reaction mentioned above takes place, and the carbon 
monoxide and hydrogen thus formed pass into a large, 
vertical iron cylinder filled with brick checker-work. In 
this cylinder, which has been previously heated, the gases 
come in contact with hydrocarbon vapors obtained from 
petroleum. The gases and vapors pass on into another 
similar highly heated cylinder, where the hydrocarbons 
are so completely decomposed and converted into gases 
that they will not again condense to oils when cooled. 
After a very few minutes the steam blast has so cooled 
the coke fire that it becomes necessary to shut off the 
steam and admit the air blast again. During this proc- 
ess the two cylinders, which are known respectively as 
the carburetter and superheater, are reheated by the 
combustion of the gas in them. Water gas thus formed 
is purified by passing through condensers and scrubbers 
and iron oxide purifiers, and is then sometimes mixed 
with coal gas and is ready for use. 



SUMMARY 

All fuels contain carbon. There are three forms of carbon — 
diamond, graphite, and amorphous carbon. The last in- 
cludes all the uncrystallized varieties, such as charcoal, coal, 
coke. 

The Diamond. — The diamond is very hard, crystallized, and when 
properly cut and polished is often very brilliant. Its specific 
gravity is 3.5. 

Graphite. — Graphite is made up of hard, slippery particles. It 



84 A PRACTICAL CHEMISTRY 

withstands high temperatures, and is used on this account 
for making crucibles. It is also used for stove polish, paint, 
to lubricate machinery, and in the making of lead pencils. 

Wood. — Wood contains cellulose and mineral matter. When 
burned the latter forms the ash, while cellulose is converted 
into water and carbon dioxide. 

Destructive Distillation. — Destructive distillation is the process 
of decomposing a substance into new substances by heating 
it with little or no air. 

Charcoal. — Charcoal is a product of the destructive distillation of 
wood. It is a black, porous substance capable of absorbing 
gases and of destroying coloring matter. It is used as a 
fuel and for filtering. 

Bone Black — Animal Charcoal. — Boneblack and animal charcoal 
are forms of charcoal obtained by the destructive distillation 
of bones and of blood mixed with sodium carbonate. They 
are used in refining sugar. 

Coal. — Coal has been formed in past ages by the decomposition 
of wood in swamps. The many kinds of coal may be divided 
into two classes — bituminous and anthracite. The combus- 
tion of coal produces ash, carbon dioxide, carbon monoxide, 
nitrogen, sulphur dioxide, and hydrogen sulphide. The car- 
bon monoxide is a dangerous poison. 

Hydrocarbons. — Hydrocarbons are compounds of carbon and 
hydrogen only. They include such substances as methane, 
which is also known as marsh gas, and fire damp, and 
acetylene. Mixtures of fire damp and air produce explosions 
in coal mines. 

Petroleum. — Petroleum is a mixture of hydrocarbons obtained 
as oil from the earth. By distillation it is separated into a 
number of oils, as gasoline, benzine, kerosene, paraffin oil, 
and other substances. 

Destructive Distillation of Coal. — The destructive distillation of 
coal produces coke (which is a hard, porous fuel producing 
much heat in combustion), ammonia, coal gas, and tar. 
These are separated by mechanical and chemical processes. 
The purified coal gas is used for illumination and other 
familiar purposes. 



CARBON AND FUELS 85 

Water Gas. — The water gas industry is founded upon the re- 
action, C + H 2 = H^ + CO. The hydrogen and carbon 
monoxide mixture is enriched with hydrocarbons and purified. 

REVIEW QUESTIONS 

1. What is a fuel? 

2. How is the study of carbon connected with the study 
of fuels? 

3. What evidence have we that all forms of carbon contain 
the same element? 

4. Upon what properties of the diamond does its use as a 
gem depend? 

5. How does graphite differ from diamond? 

6. What use is made of black diamonds? 

7. Mention some uses of graphite and tell upon which of its 
properties each use depends. 

8. Why should wood come first in a study of fuels? 

9. W x hat are the products of the combustion of wood? 

10. How do wood ashes differ from coal ashes? 

11. What are some of the other products formed when char- 
coal is made? 

12. What is destructive distillation? 

13. Name some uses of charcoal in the home. 

14. Tell some of the changes which have occurred in the earth 
from time to time, and tell how these changes are connected 
with the formation of coal. 

15. By what process is sugar made white? 

16. What are the evidences that coal is formed from wood? 

17. Which produces the more flame, bituminous or anthracite 
coal? 

18. How is carbon monoxide formed in the coal fire? 

19. Why should you not have a cracked stove in a room 
where you sleep? 

20. In what way may part of the fuel in a coal fire be lost? 

21. What is the smoke from soft coal? 

22. How does the manufacture of coke compare with that, 
of charcoal? 



86 A PRACTICAL CHEMISTRY 

23. What method of making coke would you think most 
profitable? 

24. Explain how the use of coke depends upon its properties. 

25. What are hydrocarbons? 

26. What modern invention has brought acetylene into use? 

27. Give the equation for the formation of acetylene. 

28. Explain the acetylene burner. 

29. How would you cut a steel beam? 

30. What does the name "petroleum" mean? 

31. Why do oil wells "gush" at first and later need to be 
pumped? 

32. What is the composition of American petroleum? 

33. Why is it necessary to refine kerosene? ' 

34. Describe the destructive distillation of soft coal. 

35. Explain how the various products are separated. 

36. What uses can you think of for coal tar? 

37. Does "water gas" contain water? 

38. Upon what principle does its manufacture depend? 

39. What treatment does this gas require when used for il- 
luminating ? 



CHAPTER X 
FLAMES 

In the study of the combustion of fuels it is noticed 
that some substances burn with flame and others with 
little or none. A flame is a burning gas; all substances, 
therefore, which burn with flames must be either gases or 
capable of being converted into gas by the heat which 
they produce during combustion. Fuels, like charcoal, 
from which nearly all volatile matter has been removed 
must burn almost without flame. 

There are many things of interest and value which 
may be learned by a study of the shape, structure, com- 
position, color, light, and heat of flames, and some of 
these will be discussed in this chapter. 

Structure of Flames. — If we examine such flames as 
those of the candle, Bunsen burner, and ordinary gas jet, 
we find near their middle portion a zone of unburned gas 
which has not yet been heated to the kindling tempera- 
ture. This is surrounded by a zone in which combustion 
is complete, and this in turn is nearly inclosed in an en- 
velope of products of combustion. The unburned gas 
may be shown to exist in a candle flame by putting the 
end of a glass tube near the wick and burning the gas 
as it passes out at the other end of the tube. The pres- 
ence of a zone without combustion in the middle of a 
Bunsen flame may be shown by suddenly thrusting an 
unburned match into the flame, where it may be held for 
some time without catching fire. Another way to study 

87 



A PRACTICAL CHEMISTRY 



the structure of a candle flame is to hold a sheet of paper 
down horizontally upon the flame till the paper begins to 
scorch. Again a longitudinal section may be obtained 
by holding the paper upright against the flame. 

Oxidizing and Reducing Flames. — Most flames contain 
much hot carbon in the process of oxidation and are 
therefore ready to take up oxygen from other substances 

containing it. These are 
known as reducing flames. 
Those flames, on the 
other hand, which contain 
an excess of oxygen and 
can give it to other sub- 
stances are called oxidiz- 
ing flames. Both oxidiz- 
ing and reducing flames 
can be made with the 
blowpipe, as shown in the 
illustration. The air is 
shut off from the Bunsen 
burner and the gas turned 
down till the flame is a 
little more than an inch 
high, the blowpipe tip is 
placed within the flame, 
and a blast of air blown 
with sufficient force to 
produce a blue tonguelike flame. This is the oxidizing 
flame much used in blowpipe work. To produce the blow- 
pipe reducing flame the tip of the pipe is held just out- 
side of the Bunsen burner and a gentle blast of air is 
blown with only enough force to carry the yellow cone 
of reducing flame over the object to be reduced and sur- 
round it with the burning carbon. 

On a large scale reducing flames are used in furnaces 




Fig. 34. — The Blast Lamp. The blast 
lamp is similar to the oxyhydrogen 
blowpipe in construction. Illuminating 
gas is used in the outer tube and a 
blast of air in the inner tube. It is 
much used in the laboratory when a high 
temperature is required. 



FLAMES 



89 




Fig. 35. — Reducing and Oxidizing Flames. The reducing flame, R, is made by hold- 
ing the blowpipe, B, just outside of the Bunsen flame. In making the oxidizing 
flame, 0, the blowpipe tip is within the flame and the air is blown with more force. 



to reduce metals from their ores. The acetylene torch as 
used to cut steel beams is an example of an oxidizing 
flame. 

Color, Temperature, Light. — The color of flames is often 
due to some particular element that is burning. Thus the 
vapor of sodium gives a yellow 
flame ; potassium gives violet ; 
barium; green. On the other hand, 
there are a number of compounds 
which give characteristic flames. 
A familiar example is carbon 
monoxide. 

The temperature of most flames 
depends largely upon the rapidity 
of the combustion, but is also de- 
pendent upon the character of the 
burning gas, since some fuels con- 
tain much more energy than 
others. 

Many factors are involved in the 
light-producing power of flames, 
such as temperature, pressure, ra- 
pidity of combustion, amount of 
oxygen present, and, above all, the 

presence of solid carbon particles derived from hydro- 
carbons. If a gas contains no solid particles, as, for ex- 




Fig. 36.— A Gas Mantle. 




OftCJ4S»£ 

bfl P PI £~ « 
>> o 

03 J 



o3<£ S'S oT-c 
2^uh § £ 3 






- r3 += 



£3 



*- 2 1 rs 

W 03 Pi > g £ 
°= M 2 rt 

„ c3 fl 43 
«! 

O i- 

w 03 rt C 

5 J3 ^ -S -^ 

p2 8^-2 6 
2 o ropn t, ox 

Sci £ § .sp 
sg~,§J|| 

J; C on u<g.=s 

. n ni i« naK 
g 



92 



A PRACTICAL CHEMISTRY 




Fig. 39. — A Gauze Placed Over a 
Burner with the Escaping Gas 
Lighted Above the Gauze. The 
flame cannot pass through the gauze 
because the latter conducts the heat 
away so fast that the temperature 
of the gas does not reach the kin- 
dling point. 



ample, hydrogen or carbon 
monoxide, the flame is prac- 
tically non-lnminous. Such 
flames can be rendered lumi- 
nous by putting into them 
solid matter, as lime, iron 
filings, and charcoal, or a gas 
mantle. If the gas contains 
hydrocarbons, the same re- 
sult is obtained. The hydro- 
gen of these compounds 
burns before the carbon, but 
in burning heats the carbon 
to incandescence. In the 
presence of a large quantity 
of oxygen the combustion of 
the carbon is so rapid that 

the particles do not have time to become white hot before 

they are oxidized. This is. illustrated 

by the Bunsen burner. If a cold object 

be held in the flame when the air holes 

of the burner are closed, the burning 

particles of carbon are so chilled that 

they stop burning and deposit on the 

cold surface. If the air holes are not 

closed so that the flame has plenty of 

oxygen, carbon will not be thus depos- 
ited. 

Wire Gauze and Flames. — Many 

years ago Sir Humphrey Davy invented 

the miners' safety lamp, based upon 

the theory that flame will not pass 

through fine wire gauze because the wire 

conducts the heat away so fast that the 

p . , , . . Fig. 40. — Davy's Safe- 

temperature of the burning gas is re- TT Lamp> 




FLAMES 



93 



duced below the kindling point. This may be illustrated 
by holding a gauze above a burner and lighting the escap- 
ing gas above the gauze. The gas will burn only above 
the gauze. 

In Davy's lamp the flame was sur- 
rounded by fine wire gauze, which pre- 
vented the fire damp in the mine from 
catching fire and exploding. Inside of 
the wire gauze the mixture of fire damp 
and air might burn without harm, but 
as soon as it reached the wire gauze it 
would be cooled below the kindling tem- 
perature and the flame prevented from 
passing to the outside. 

Fire Extinguishers. — Large fires are 
usually extinguished with water, but 
small ones may often be put out by 
means of chemical extinguishers such as 
are usually seen around public build- 
ings. 

In the bottom of these extinguishers is 
placed a mixture of sodium bicarbonate 
and water, while a loosely stoppered bot- 
tle containing sulphuric acid is held in a 
little frame near the top. When the ex- 
tinguisher is taken to the fire and in- 
verted the stopper falls from the acid 
bottle, the acid flows out and mingles 
with the sodium bicarbonate, liberating carbon dioxide. 
The carbon dioxide surrounds the fire and excludes the 
oxygen of the air. The formation of the carbon dioxide 
may be represented by the equation : 2NaHC0 3 + H 2 S0 4 
- Na 2 S0 4 + C0 2 + 2H 2 0. 




Fig. 41. — Fire Extin- 
guisher. C is the 
solution of carbo- 
nate; A, the bottle 
of acid with its 
loose fitting stopper 
S; H is a short hose, 
through which the 
contents of the ex- 
tinguisher are 
squirted upon the 
fire. For recharging 
the tank, the top, 
T, is unscrewed. 



94 A PRACTICAL CHEMISTRY 

SUMMARY 

A flame is a burning gas. Flames may be studied as to shape, 
structure, composition, color, light, and heat-giving power. 

The middle zone of a flame is composed of unburned gas. This 
is surrounded by a zone of combustion which is nearly in- 
closed by the products of combustion. 

Reducing Flames. — Reducing flames contain much carbon and 
hydrogen and remove oxygen from other substances. 

Oxidizing Flames. — Oxidizing flames contain oxygen with which 
they can oxidize other substances. 

The color of flames is due to the nature of the burning substance, 
while their temperature depends largely upon the rapidity of 
the combustion. 

Flames are rendered luminous by the presence of solid matter, 
usually particles of carbon derived from the decomposition 
of hydrocarbons in the gas. 

The use of the Davy safety lamp depends upon the fact 
that flames do not pass through fine wire gauze. In the 
chemical fire extinguisher carbon dioxide is formed by the 
decomposition of a carbonate by sulphuric acid. The carbon 
dioxide surrounds the fire and excludes the air. 



REVIEW QUESTIONS 

1. What is a flame? 

2. Why does burning charcoal produce but little flame? 

3. What are some of the subjects of study in connection 
with flames? 

4. Why is it difficult to extinguish burning gasoline? 

5. Which part of a burner flame is the hottest? 
C. What are oxidizing and reducing flames? 

7. Is any part of a Bunsen flame a reducing flame? 

8. Give some examples of the uses of reducing and oxidizing 
flames. 

9. Can you suggest a method of determining what elements 
are in the stars? 



FLAMES 95 

10. Why does not the gas mantle burn up? 

11. Explain the limelight. 

12. Explain the principle of Davy's safety lamp. 

13. How can you illustrate this"? 

14. Explain the action of the chemical fire extinguishers. 



CHAPTER XI 
PHOSPHORUS AND MATCHES 

The last two chapters have been devoted to fuels and 
flames. The means of starting combustion now claim 
our thought. The methods of the primitive man for ob- 
taining fire were very crude and accomplished only with 
patience and much hard work. Progress in improving 
these methods was slow, and even in the memory of men 
now living the flint and steel and flintlock gun were in 
common use. 

All the numerous varieties of our modern friction 
match require phosphorus in some form for their con- 
struction; hence it is desirable that we study this im- 
portant element before we try to understand their manu- 
facture. 

Discovery. — It was about the year 1666 that Brandt 
first obtained phosphorus from a crystalline substance 
known as microcosmic salt. The chemistry of the ele- 
ment has been worked out little by little since that time. 

Sources. — Our modern chemical name for microcosmic 
salt is hydrogen sodium ammonium phosphate. It has 
the formula, HNaNHVP0 4 . This means it belongs to that 
large class of substances called phosphates, a number of 
which occur in rocks and soils. We also find phosphates 
in bones and plants, and other phosphorus compounds in 
animal tissues. In South Carolina, Florida, Tennessee, 
and other parts of the world are large deposits of phos- 
phate rock which contain much calcium phosphate, 



PHOSPHORUS AND MATCHES 



97 



Ca 3 (P0 4 ) 2 . From these various sources the world's sup- 
ply of phosphorus and phosphorus compounds is ob- 
tained. 

Manufacture. — The modern method for obtaining phos- 
phorus is by means of an electric furnace somewhat in 
the shape of a large box containing electrodes placed 
horizontally. The furnace has an outlet at the bottom 




Fig. 42. — Phosphorus Furnace. The mixture of phosphate, carbon and silicon 
dioxide is fed into the hopper, H, with the slide, T, open. T is then pushed in 
and T/ drawn out. The screw, S, carries the mixture and drops it between the 
electrodes, A and C. The slag which is formed flows out at O, while the gases pass 
through R. The phosphorus vapor condenses in the water at P and the oxide of 
carbon bubbles out at G. 



for the slag and another near the top for the phosphorus 
vapor. A mixture of calcium phosphate, sand, and car- 
bon is admitted at the top and allowed to drop down be- 
tween the electrodes through the electric arc. The great 
heat of the electric arc brings about a reaction between 
the calcium phosphate, sand and carbon, forming calcium 
silicate of the sand and calcium, which becomes a slag 
and flows out at the bottom. The carbon and oxygen 



98 A PRACTICAL CHEMISTRY 

form carbon monoxide or dioxide, while the phosphorus 
passes out as vapor and is condensed underwater. The 
explanation of these reactions is found in the facts: (1) 
That sand, which is silicon dioxide, Si0 2 , when fused with 
calcium phosphate, forms calcium silicate, CaSi0 3 , and 
phosphorus pentoxide, P 2 5 , as in the equation, 

Ca 3 (P0 4 ) 2 + 3Si0 2 = 3CaSi0 3 + P 2 5 - 

(2) That phosphorus pentoxide can be reduced when 
heated to a high temperature with carbon, 

P 2 5 + 5C = 2P + 5CO. . 
Carbon 
monoxide. 

Forms. — After having learned of two forms of oxygen 
and three of carbon, we are not surprised to find that 
phosphorus exists in several modifications. The most 
active form is usually called waxy phosphorus, but it is 
sometimes spoken of as yellow, and sometimes as white, 
phosphorus. At the ordinary temperature it is a yellow, 
waxy substance which becomes quite hard and brittle 
when cold. It melts under water at 44° C, and will take 
fire spontaneously when exposed to the air. On this 
account it must be kept under water. "When exposed 
to the light it sometimes becomes coated with a white 
substance. Occasionally this coating is a bright red. 
Waxy phosphorus is soluble in carbon bisulphide. It is 
extremely poisonous when taken internally, and its vapor 
when inhaled produces a rotting of the bones. 

When ordinary waxy phosphorus is heated out of con- 
tact with the air to a temperature of about 275° C. it is 
changed into a dark red brown powder called red phos- 
phorus. This is very different in properties from waxy 
phosphorus; it is insoluble in carbon bisulphide, will not 
burn easily, is not volatile or poisonous. At a higher 



PHOSPHORUS AND MATCHES 99 

temperature it is converted again into waxy phosphorus. 

One or two crystalline forms of phosphorus can be 
made, but as yet are of but little practical importance 
and for our purposes need not be studied. 

Compounds. — The compounds of phosphorus are quite 
numerous. When burned it unites with oxygen and 
forms two oxides known as phosphorus trioxide and 
phosphorus pentoxide. The latter is the more important, 
and is the chief product formed when the element is 
burned in a liberal supply of air or oxygen. The trioxide 
is formed when the supply of oxygen is limited. Phos- 
phorus pentoxide is a white solid which dissolves in 
water, forming phosphoric acid. (Note. — An oxide 
which forms an acid on dissolving in water is called an 
anhydride.) Several different phosphoric acids may be 
formed, depending upon the number of molecules of 
water reacting with one molecule of phosphorus pent- 
oxide. 

The various phosphates are the salts of these phos- 
phoric acids. Some of them are used in fertilizers, and 
we shall have occasion to take them up again under that 
heading. 

With hydrogen, phosphorus forms several compounds 
called phosphenes, which are interesting on account of 
the ease with which they burn. The most common of 
these phosphenes, when freshly made, burns spon- 
taneously when exposed to the air. 

Matches. — The most important use of phosphorus is 
in the manufacture of matches, which are used by the 
millions and are made in a number of different ways. 

The process of lighting a match consists in firing by 
means of friction a small amount of phosphorus, or a 
sulphide of phosphorus, which in turn sets fire to some 
easily combustible substance mixed with an oxidizing 
agent, and this again burns the stick. 



100 A PRACTICAL CHEMISTRY. 

In the ordinary match a small bit of waxy phosphorus 
mixed with a little glue, some oxidizing agent,, and 
coloring matter is stuck onto the end of the stick. Sur- 
rounding the sides of this but leaving the top exposed, 
there is a band of paraffin, sulphur, or some other com- 
bustible substance mixed with an oxidizing agent. For 
several reasons these matches are objectionable; they 
are often ignited by being stepped upon; sometimes 
when lighted the match head flies to pieces and sets fire 
to clothing and other inflammable^ material, and they 
often produce unpleasant odors. 

Safety matches are constructed on a different plan. A 
mixture of inflammable material, such as antimony sul- 
phide, manganese dioxide and glue, is put onto the match 
stick, while on the box are red phosphorus, powdered 
glass and glue. These matches are free from some of 
the objectionable features of the ordinary match, but 
are not so convenient to use, since they must be struck 
upon the box. 

On account of the poisonous vapor of ordinary phos- 
phorus, the workmen in match factories often suffer 
dreadfully. In recent years much effort has been made 
to substitute phosphorus sesquisulphide, a harmless sub- 
stance, for phosphorus and thus protect the workmen. 



SUMMARY 

Phosphorus. — Phosphorus was. discovered by Brandt about the 
year 1666. It is found in rocks and plants and in the bones 
and other tissues of animals. 

Phosphorus is obtained by heating calcium phosphate mixed with 
sand and carbon in an electric furnace. 

Phosphorus has several modifications, as waxy or yellow phos- 
phorus, and red phosphorus. These may be converted into 
each other, but are quite different in properties. 



PHOSPHORUS AND MATCHES 101 

Many compounds of this element, including oxides and acids, are 
known. 

On account of its low kindling point, waxy phosphorus is much 
used in making matches. Red phosphorus is used in mak- 
ing safety matches. Phosphorus sesquisulphide is also 
used for the same purpose in place of waxy phosphorus. 
Its use protects the workmen in the match factory from 
phosphorus poisoning. 



REVIEW QUESTIONS 

1. In what respect does the ancient method of kindling fires 
resemble the modern? 

2. Phosphorus means light bearer; how does it get this 
name? 

3. What parts of the body contain the most phosphorus? 

4. Describe the modern method for obtaining phosphorus. 

5. Explain the chemistry of the process. 

6. Compare the properties of red phosphorus with those of 
waxy phosphorus. 

7. Name two oxides of phosphorus and tell how they are 
made. 

8. Explain the relation of phosphorus pentoxide to the 
phosphoric acids. 

9. What products are formed when hydrogen compounds of 
phosphorus burn? 

10. Explain the lighting of a match. 

11. What is the composition of most ordinary matches? 

12. What would be the result if the oxidizing agent were put 
on the box in place of on the stick of the safety match? 

13. How would the use of phosphorus sesquisulphide improve 
the condition of the workmen in match factories? 



CHAPTER XII 
SULPHUR AND SOME OF ITS COMPOUNDS 

On account of the great number of uses for which 
sulphur and its compounds are employed, it is best that 
some time be given to its study before going on to 
other things. 

Sulphur is found in many parts of the earth, both in 
the free state and in compounds. The free sulphur is 
mostly found in the vicinity of volcanoes, as in Sicily, 
where until recent years much of our sulphur was ob- 
tained. Lately, however, this country has produced its 
own sulphur. 

The free, or native, sulphur is often found in crystal- 
line form, some of the crystals being of considerable 
size and beauty and of a yellow color. Much of it, how- 
ever, is in large masses mixed with earthy impurities. 
To separate the sulphur from these impurities it is first 
melted and allowed to run through a grate away from 
the earthy matter; after this it is distilled from a retort 
and the vapor condensed in brick chambers. While the 
chambers are cool the sulphur collects on the walls as 
fine yellow crystals mixed with a white amorphous pow- 
der. This mixture is called flowers of sulphur ; but when 
the walls are heated the crystals can no longer form, and 
the sulphur runs to the bottom of the room as a' liquid, 
which may be cast into rolls in wooden molds. In this 
form it is called roll sulphur. This process of purification 
is employed with the sulphur obtained in Sicily. In Lou- 

102 






SULPHUR AND SOME OF ITS COMPOUNDS 103 

isiana the sulphur deposit is from 800 to 1,000 feet below 
the surface of the ground. Wells are drilled into this de- 
posit and several pipes, one within the other, are driven 
into each well. Water, heated under pressure to a tem- 
perature of 168° C, is forced down into the sulphur 
through one of the pipes. This superheated water melts 
the sulphur and causes it to rise in another of the pipes. 




Fig. 43. — Sulphur Flowing from a Sulphur Well. The sulphur on being forced 
from the ground in the melted condition flows into a large vat, where it soon be- 
comes solid. Courtesy of the Union Sulphur Company. 

The pressure of the water, however, is not sufficient to 
lift the melted sulphur to the top of the ground and must 
be aided by compressed air forced down through a third 
pipe. After coming to the surface the sulphur runs into 
large wooden tanks, as shown in Figure 43. The sulphur 
thus obtained is nearly pure. After the sulphur solidifies 
the tanks are torn away, leaving solid blocks of sulphur 
of many thousands of tons weight. (See Figure 44.) 



104 



A PRACTICAL CHEMISTRY 



These are blasted to pieces and are ready for shipment. 
For some purposes this sulphur must be purified. 

Sulphur crystallizes in two forms. It may also be con- 
verted into a very interesting amorphous condition by 
heating nearly to boiling and then suddenly cooling. 
This is called plastic sulphur. "When first made it is 
quite elastic, but soon becomes brittle on standing. 




Fig. 44. — Many Thousands op Tons op Sulphur in One Pile. 



The numerous changes which sulphur passes through 
when heated are very interesting. At first it melts to 
a thin, lemon yellow liquid, which quickly changes to 
amber, and finally to such a dark red brown that it seems 
almost black. At the same time the liquid becomes 
thicker and more viscous until it is jelly-like. After this, 
if the heating is continued, it becomes thinner again and 
begins to boil. 

Oxides of Sulphur. — Two of the several oxides which 



SULPHUR AND SOME OF ITS COMPOUNDS 105 

sulphur forms are of much value iu manufacturing. The 
first of these, sulphur dioxide, S0 2 , is made by burning 
sulphur. Both the free sulphur and sulphides are used 
for this purpose. Iron pyrites (FeS 2 ), a mineral com- 
posed of iron and sulphur, is the sulphide most used. 
This is burned like coal, the iron forming ferric oxide 
and remaining on the grate, while the sulphur dioxide, 
which is a colorless, suffocating, irritating gas, is con- 
ducted away through a flue to the place where it is used. 
The reaction is represented thus : 

4FeS 2 + 110 2 = 8S0 2 + 2Fe 2 3 

Sulphur dioxide will not support combustion, nor will it 
burn., It dissolves in water, forming a solution which 
smells of the oxide. This solution has acid properties 
and contains sulphurous acid (H 2 S0 3 ). With bases it 
forms salts called sulphites. These are of great value in 
paper making. 

Sulphur dioxide is used as a disinfectant to kill the 
germs of contagious diseases, as a preservative, and as 
a bleaching agent. These uses will be considered again. 
Its greatest use, however, is in the manufacture of sul- 
phuric acid. 

Sulphuric Acid. — In the manufacture of sulphuric acid 
sulphur dioxide is used in two different ways, depending 
on the methods of manufacture employed. These are 
called the contact method and the lead chamber process. 

In the contact method sulphur dioxide, which has been 
obtained by the burning of iron pyrites, is converted 
into sulphur trioxide by heating the dioxide and air in 
contact with a porous mass called a catalytic agent or 
a catalyzer, as represented in the equation, 2S0 2 + 2 = 
2S0 3 The action of catalyzers is not yet thoroughly un- 
derstood. They are substances which will hasten or re- 
tard a chemical change. The use of manganese dioxide 



106 



A PRACTICAL CHEMISTRY 



with potassium chlorate in the preparation of oxygen 
(page 20) is an illustration of a .catalyzer. Finely 
divided platinum is a good catalyzer for converting sul- 
phur dioxide into the trioxide. The sulphur trioxide is 
next dissolved in water, forming sulphuric acid. Thus, 
S0 3 + H 2 = H 2 S0 4 . Here again a catalytic agent is 
found desirable. This time sulphuric acid itself is the 
catalyzer. The sulphur trioxide is passed into water 
containing sulphuric acid until the desired concentration 
is obtained. The sulphur dioxide, if made from iron 

Lead Chambers 




To 
Chimney 




Fig. 45. — Diagram of a Sulphuric Acid Plant. 



pyrites, must be purified by washing and filtering before 
it is made into sulphur trioxide. 

In the other method of manufacturing sulphuric acid 
a very different process is followed, called the chamber 
process. Three steps are involved in this method: (1) 
The making of the sulphur dioxide, which is accomplished 
in the same ways as in the contact process. (2) The for- 
mation of sulphurous acid by the union of water with 
the sulphur dioxide. (3) The oxidation of the sulphur- 
ous acid to sulphuric by means of oxides of nitrogen and 
the oxygen of the air. The second and third steps are 
accomplished at the same time. The process is described 
in connection with Figures 45, 46, and 47. Figure 45 is a 



SULPHUR AND SOME OF ITS COMPOUNDS 107 



Pyrite 
Furnaces 



fcQ 



-*-To 
First Lead, 
Chamber 



diagram of a sulphuric acid plant. In the pyrite burners, 
iron pyrites is burned as coal is burned in other furnaces 
and sulphur dioxide is formed as a product. The sulphur 
dioxide mixed with air passes into the Glover tower, 
which is shown in Figure 46. Here it is mixed with oxides 
of nitrogen and steam and the whole mixture passes into 
the lead chambers. In these chambers more steam is 
added and the chemical changes occur which produce sul- 
phuric acid. The mixture of gases which passes out of 
the last of the lead chambers con- 
sists mostly of nitrogen from the air 
and of oxides of nitrogen. In order 
to separate and save the oxides of 
nitrogen, this mixture of gases is 
forced up through the coke-filled 
Gay-Lussac tower shown in Figure 
47. Concentrated sulphuric acid 
trickling down through this tower 
from a tank at the top meets the 
mixture of gases and absorbs the 
oxides of nitrogen while the remain- 
ing gases escape to the chimney. 
The solution of oxides of nitrogen 
in concentrated sulphuric acid is 
pumped to the top of the Glover 

tower. Here it is mixed with water and caused to trickle 
down over stones in this tower. On the way down, the 
oxides of nitrogen escape from the acid and again enter the 
lead chambers to be used in oxidizing more sulphurous acid 
into sulphuric acid, for none of these oxides have been 
destroyed. They serve as carriers of oxygen from the air 
to the sulphuric acid. Some of the oxides of nitrogen, 
however, are lost in the operation, and to make good this 
loss more nitric acid must, from time to time, be sent into 
the chambers. This is done by heating a mixture of so- 



Fig. 46. — The Glover Tow- 
er in a Sulphuric Acid 
Plant. 



108 



A PRACTICAL CHEMISTRY 



umneyji 



m 






sag 



dium nitrate and sulphuric acid and passing the nitric 
acid vapors and oxides of nitrogen thus formed into the 
leaden chambers. The sulphuric acid, when taken from 
the chambers, contains much water, and must be concen- 
trated by boiling off the water first in leaden pans and 
afterward in cast iron or platinum pans. The choice 
of a metal for the pans depends upon the fact that lead 
is not very soluble in dilute sulphuric acid, but quite 
soluble in concentrated sulphuric 
acid, while cast iron is easily soluble 
in dilute sulphuric acid and nearly 
insoluble in the concentrated acid. 

Sulphuric acid is a heavy, oily, 
colorless liquid, a little more than 1.8 
times as heavy as water. Its boiling 
point is quite high, and until near its 
boiling point it has no odor. The 
strong acid rapidly absorbs water 
from the air and from other things 
and will decompose most organic mat- 
ter, abstracting from it the elements 
of water. The carbon of the organic 
matter at the same time turns the acid 
black. On the other hand, hot reduc- 
ing agents, such as carbon, copper, 

sulphur, and organic matter, decompose sulphuric acid with 
the formation of sulphur dioxide, water and oxides of the 
reducing agents. This fact is made use of as a laboratory 
method of generating sulphur dioxide. Sulphuric acid and 
copper are heated together and sulphur dioxide is given 
off. In the laboratory sulphur dioxide may also be formed 
by decomposing a sulphite with an acid. The following 
reactions illustrate these two methods of generating sulphur 
dioxide : 



Qt 



„ rom 

B£ Fourth 

Lead 

Chamber 



Fig. 47. — The Gay-Lus- 
sac Tower in a Sul- 
phuric Ach> Plant. 



SULPHUR AND SOME OF ITS COMPOUNDS 109 
(l).With reducing agents, 

C + 2H 2 S0 4 = C0 2 + 2H 2 S0 3 

(Sulphurous 
acid) 

and heat decomposes 

H 2 S0 3 =:H 2 + S0 2 

or, combining these equations : 

C + 2H 2 S0 4 = C0 2 + 2S0 2 + 2H 2 

(2) Decomposition of a sulphite, 

Na 2 S0 3 + 2HC1 = 2NaCl + H 2 S0 3 
(Sodium 
sulphite) 

and H 2 S0 3 = H 2 + S0 2 , 

or, combining the two equations, 

Na 2 S0 3 + 2HC1 = 2NaCl + H 2 + S0 2 . 

Hydrosulphuric Acid and Sulphides. — Corresponding 
to hydrochloric acid, we have another sulphur acid, hy- 
drosulphuric acid, commonly called hydrogen sulphide. 
It consists of hydrogen and sulphur, and is a poisonous 
gas with an unpleasant odor. It is soluble in water, and, 
while not a strong acid, forms many salts called sul- 
phides. These sulphides may also be made by the direct 
union of metals with sulphur. Hydrogen sulphide is 
usually made by the action of a dilute acid on a sulphide, 
particularly ferrous sulphide. Thus, 

FeS + H 2 S0 4 = FeS0 4 + H 2 S. 
Ferrous Ferrous Hydrosulphuric 

sulphide sulphate acid 



110 A PRACTICAL CHEMISTRY 

Sulphur also forms sulphides with some of the non- 
metallic elements like carbon and phosphorus. The sul- 
phide of carbon is called carbon bisulphide. It is a vola- 
tile, inflammable liquid, the vapors of which are poison- 
ous. In the impure form, as it is usually sold, it has a 
very unpleasant odor. It is made by passing sulphur 
vapor over heated carbon. Its formula is GS 2 The mod- 
ern plan is to keep the carbon heated and the sulphur 
vaporized by means of an electric arc near the bottom 
of the apparatus. The carbon bisulphide (or disulphide, 
as it is sometimes called) passes out as a vapor and is 
condensed. Carbon bisulphide is used as a solvent for 
rubber, fats, and gums. Sulphur is also quite soluble in 
carbon bisulphide, and from this solution the sulphur 
crystallizes. Phosphorus also dissolves in this liquid. 



SUMMARY 

Sulphur. — Sulphur occurs both free and combined in the earth. 
Free sulphur has many crystalline forms, and may also be 
converted into the amorphous condition. It passes through 
many changes of color and density during heating. 

Sulphur Dioxide. — Sulphur dioxide is the product of the com- 
bustion of sulphur and sulphides in the air. It forms sul- 
phurous acid with water. It is much used in making sul- 
phur trioxide and sulphuric acid. Sulphur dioxide is also 
used as a disinfectant and in the manufacture of sulphites. 
It can be oxidized in the presence of a catalytic agent. 

Sulphur Trioxide. — Sulphur trioxide is made from the dioxide by 
heating the latter with oxygen in the presence of a catalyzer. 
Sulphur trioxide with water forms sulphuric acid. 

Sulphuric Acid. — Sulphuric acid is also made by the chamber 
process, which consists in the making of sulphur dioxide by 
burning iron pyrites, the uniting of the dioxide with water, 
and the oxidation of the product by means of oxides of 
nitrogen and the oxygen of the air. The dilute acid thus 



SULPHUR AND SOME OF ITS COMPOUNDS 111 

formed must be concentrated by evaporation. This is done 
first in lead, and afterward in cast iron pans. 

Sulphuric acid is a heavy, colorless liquid which absorbs 
water from the air and decomposes organic matter, liberat- 
ing carbon, which gives it a dark color. It is decomposed 
by hot reducing agents. 
Sulphides. — Sulphides are salts of hydrosulphuric acid, but many 
of them may be made by direct union of sulphur with 
metals. Hydrosulphuric acid is usually made by the action 
of an acid upon a sulphide. It is a poisonous gas. Some 
useful sulphides of non-metallic elements are known. 



REVIEW QUESTIONS 

1. Name a mineral containing combined sulphur. 

2. What is meant by native sulphur? 

3. By what three methods can sulphur be crystallized 1 ? 

4. Although workmen receive more pay in this country than 
in Sicily, the cost per ton of producing sulphur in Sicily is 
greater than in Louisiana; explain why. 

5. Describe the changes in sulphur during heating. 

6. Which do you think should be the more expensive and 
why, flowers of sulphur or roll sulphur"? 

7. If zinc sulphide were burned what products would be 
formed ? 

8. What are the properties of sulphur dioxide? 

9. Does a solution of S0 3 in water smell? 

10. How can you make sulphites? 

11. How does a solution of S0 2 preserve fruits and vege- 
tables? 

12. What is a catalytic agent? 

13. Describe the contact method of manufacture of sulphuric 
acid. 

14. How does the chamber process differ from the contact 
method? 

15. Explain how the oxides of nitrogen are saved. 

16. Why are the chambers lined with lead in place of wood? 



112 A PRACTICAL CHEMISTRY 

17. How is the supply of oxides of nitrogen replenished? 

18. What is dilute sulphuric acid? 

19. Give the properties of sulphuric acid. 

20. Explain the laboratory methods for making sulphur 
dioxide. 

21. Can you burn hydrogen sulphide? 

22. How would you make ferrous sulphide? 

23. What products are formed when ferrous sulphide reacts 
with dilute sulphuric acid? 

24. What products are formed by the combustion of carbon 
bisulphide ? 

25. How does carbon bisulphide differ from metallic sulphides ? 



CHAPTER XIII 
COMPOUNDS OF NITROGEN AND EXPLOSIVES 

Nitrogen as it exists in the air is an inactive element, 
uniting with great difficulty with only a few elements 
and compounds. It burns in the electric arc, forming 
oxides. At high temperatures it unites with a few metals, 
such as magnesium and calcium, forming nitrides of these 
metals. Inactive as free nitrogen may seem to be, when 
once brought into combination few elements are more 
active. The ease and rapidity with which it goes through 
various chemical changes, passing from compound to 
compound and into the free state, increase its importance 
and make it one of the most useful elements. And while 
the number of compounds of nitrogen is quite large, prac- 
tically all of them can be derived directly or indirectly 
from ammonia and nitric acid. These two compounds, 
therefore, shall first receive our consideration. 

Nitric Acid (HN0 3 ). — This substance probably does 
not occur free in nature, but several of its salts do. These 
salts are called nitrates. Sodium nitrate is found in great 
quantities in Chili. Calcium nitrate and sometimes other 
nitrates are found in the soil, while natural waters are 
seldom free from these compounds. The existence of 
nitrates in the soil is due to two causes, the combustion 
of nitrogen in the lightning flash during thunder storms 
and the action of bacteria working upon nitrogenous mat- 
ter in the soil. In the first case the products of the 
combustion are oxides of nitrogen, which with water 

113 



114 



A PRACTICAL CHEMISTRY 



change into nitric and nitrons acids, and the latter in 
turn oxidizes to nitric acid. A much larger quantity of 
nitrates, no doubt, results from the second cause, i. e., the 
action of bacteria on nitrogenous matter. In either case 
lime or other bases in the soil neutralize the acid, forming 
nitrates. 

From these nitrates, particularly from sodium nitrate, 
we obtain nitric acid by .the action of sulphuric acid. 




Fig. 48. — Making Nitric Acid. The mixture of sodium nitrate and sulphuric acid 
is heated in the retort, R. The nitric acid distills off and is condensed in the 
receiver, A, which is kept cool by a pan of water. 



The nitrate and sulphuric acid are mixed and heated in a 
retort. The nitric acid, being much more volatile than 
the sulphuric, distills off and is condensed in suitable re- 
ceivers cooled with running water. Acid sodium sul- 
phate remains in the retort. The laboratory method is 
shown in the accompanying illustration (Fig. 48). In 
the factory the acid is made on the same plan, but on a 
larger scale and with iron retorts. The receivers are of 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 115 

various forms and are made of earthenware. The chemi- 
cal reaction may be represented by the equation, 

NaN0 3 + H 2 S0 4 = NaHS0 4 + HN0 3 . 
Sodium Acid Nitric 

nitrate sodium acid 

sulphate 

Acid sodium sulphate is an acid salt. This means it is 
a compound which is both salt and acid, or, in other 
words, it is a salt in which a metal has taken the place 
of only part of the hydrogen of the acid. 

Nitric acid thus formed is a yellow liquid, the color 
being due to oxides of nitrogen dissolved in the acid. 
These oxides were formed by the decomposition of some 
of the nitric acid by the heat, and may be removed by 
blowing air through the acid. Pure nitric acid is a color- 
less liquid which boils at 86° C. Its density is 1.56. On 
standing in the light nitric acid slowly undergoes decom- 
position with the formation of oxides of nitrogen, water, 
and oxygen. In a similar way it decomposes rapidly in 
contact with any substance to which it can give its oxy- 
gen. This is well illustrated by its action upon metals. 
The metals are oxidized by the decomposing acid, and 
water and usually nitric oxide are formed. In most cases 
the metallic oxide thus formed dissolves in more of the 
nitric acid, forming water and a nitrate of the metal. 
Such metals as antimony and tin, however, form insolu- 
ble oxidation products. The oxidizing power of nitric 
acid may be shown by allowing the hot vapor to come in 
contact with wool, which it will soon ignite. On the skin 
nitric acid makes yellow stains, which can be removed 
only by removing the skin itself. Nitric acid is a very 
strong acid, as well as a strong oxidizing agent, and will 
dissolve most metals. Its salts are called nitrates. They 
are soluble in water. 



116 A PRACTICAL CHEMISTRY 

Ammonia (NH 3 ). — A very different compound of nitro- 
gen is ammonia. It is composed of hydrogen and ni- 
trogen in about the proportion of 17.6 per cent hydrogen 
to 82.4 per cent nitrogen, and is a gas nearly .6 as heavy 
as air. It has a sharp, penetrating odor and is very irri- 
tating to the nose and eyes. Ammonia is very soluble in 
water ; at 0° C, and under atmospheric pressure one liter 
of water will dissolve 1,148 liters of ammonia gas. The 
ammonia may be driven out of the water again by heat. 
The solution contains ammonium hydroxide and is called 
ammonia water, aqua ammonia, ammonium hydroxide, 
or just ammonia. It reacts in nearly all cases like the 
alkalies, sodium and potassium hydroxide, turning red 
litmus blue and neutralizing acids. Ammonia gas also 
neutralizes acids by uniting directly with them. In either 
case the product is the ammonium salt of the acid. Thus 
when hydrochloric acid and ammonia are brought to- 
gether a white cloud of ammonium chloride is formed. 

Thus, NH 3 + HC1 = NH 4 C1. 

Ammonium 
chloride 

This reaction will not take place if the gases are per- 
fectly dry. 

A word must be said about ammonium compounds: 
these compounds contain the ammonium radical (NH 4 ). 
By a radical we mean a group of elements forming part 
of a molecule and capable of being passed from one com- 
pound to another, but incapable of existing alone. The 
ammonium radical contains hydrogen and nitrogen, the 
proportion of hydrogen being greater than in ammonia. 
This radical has not been produced alone, but it has been 
obtained dissolved in mercury, the solution being known 
as ammonium amalgam. An amalgam is a solution of a 
metal in mercury. We see, then, that the ammonium 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 117 

radical is like a metal — it will dissolve in mercury, it 
will take the place of hydrogen in acids, forming salts 
called the ammonium salts, and will form a hydroxide 
(ammonium hydroxide) which behaves like the hydrox- 
ides of the alkali metals in most respects. 

Ammonia occurs in nature as the result of the action 
of bacteria on organic matter containing nitrogen ; hence 
ammonia may be found wherever such matter is de- 
caying. 

The ammonia of commerce is obtained from the de- 
structive distillation of coal. This process has already 
been described (page 80) in connection with the manu- 
facture of illuminating gas. In the laboratory ammonia 
gas is obtained by the action of slacked lime, which is 
calcium hydroxide, on ammonium chloride or some other 
ammonium salt. The mixture is heated and the calcium 
and ammonium change places, calcium chloride and am- 
monium hydroxide being formed. The ammonium hy- 
droxide is decomposed by the heat forming ammonia and 
water. In place of the slacked lime sodium or potassium 
hydroxide may be used. The equations representing the 
above changes are as follows : 

2NH 4 C1 + Ca(OH) 2 = CaCl 2 + 2NH 4 OH 
Ammonium Calcium Calcium Ammonium 
chloride hydroxide chloride hydroxide 

and NH 4 OH = NH 3 + H 2 0. 

Under ordinary conditions ammonia does not burn, but 
when mixed with oxygen it burns freely with a yellow 
flame. The products of this combustion are water and 
nitrogen. 

Uses. — Ammonia finds many uses. In solution as am- 
monium hydroxide it serves in the household as a clean- 
ing agent. It is also employed in many kinds of chemi- 



118 



A PRACTICAL CHEMISTRY 



cal work. The most important use of ammonia, how- 
ever, is in the manufacture of artificial ice. In this proc- 
ess dry ammonia gas is reduced to the liquid state by 
compressing it in pipes cooled with running water. The 
liquid ammonia next passes into coils of pipe in tanks of 
brine. In these tanks are cans of pure water to be frozen. 

JSUl 




Fig. 49. — Making Ice. I, I, the cans of water to be frozen, are in the tank of brine 
B, surrounded by the expansion coils, E, E. W is a water pipe by which the com- 
pression coils, L, are cooled. When the valve, V, is opened and V closed the 
ammonia gas in E is compressed by the pump, P, and liquefied in L. V is then 
opened and the ammonia rushes into E. 



In the coils of pipe the liquid ammonia under reduced 
pressure passes again to the gaseous state and is again 
pumped into the condensing pipes. In order that a liquid 
may change into a gas it must take up much heat, and 
the liquid ammonia while changing to gaseous ammonia 
in the coils of pipe takes heat from the brine and cans 
of water and thus cools the water till it is frozen. When 
this gas is again condensed it gives up its heat to the 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 119 

water which flows over the condenser pipes. The brine 
does not freeze, since its freezing point is much below 
that of pure water. In many cold storage plants the 
rooms are cooled by circulating the cold brine through 
pipes. Ice making is illustrated in Figure 49. 

Nitrous Acid. — A second acid of nitrogen containing 
less oxygen than nitric acid would, according to our rules 
for naming these compounds, be called nitrous acid 
(HN0 2 ). Nitrous acid is not nearly so important as 
nitric, and in fact is not even made in the pure state free 
from water. It is a much less stable compound than 
nitric acid and is not nearly so strong. Its most impor- 
tant salt, sodium nitrite, is made by taking part of the 
oxygen from sodium nitrate, Chili saltpeter. 

Sodium nitrate and lead are melted in separate ves- 
sels, then mixed and heated for several hours with con- 
stant stirring. The lead removes oxygen from the so- 
dium nitrate, forming lead oxide and sodium nitrite. 

NaN0 3 + Pb = PbO + NaN0 2 

Lead Lead Sodium 

oxide nitrite 

The mixture is then run out on a stone to cool. It is then 
broken into lumps, thrown upon a filter, and the sodium 
nitrite dissolved in water. When this solution is evap- 
orated the nitrite crystallizes. 

Sodium nitrite is used in large quantities in the manu- 
facture of a number of organic compounds, particularly 
those used as dyes. 

Nitrous acid may be made by distilling sodium nitrite 
mixed with sulphuric acid, just as nitric acid is made 
from sodium nitrate, but, as stated before, the nitrous 
acid formed in this way is not pure and free from water. 
It decomposes with heat. Strong oxidizing agents change 
it into nitric acid. During the formation of nitrates by 



120 A PRACTICAL CHEMISTRY 

the action of bacteria on decaying organic matter nitrites 
are formed as an intermediate product. Hence, the pres- 
ence of nitrites in drinking water is always regarded as 
evidence that the water is contaminated by decaying or- 
ganic matter. 

Oxides of Nitrogen. — Oxygen and nitrogen unite in 
several different proportions, forming the series of oxides 
mentioned on page 41. They are : 

Nitrous oxide, or nitrogen monoxide, N 2 0; 
Nitric oxide, NO; 

Nitrogen peroxide, N0 2 ; 

Nitrogen trioxide, N 2 3 ; 

Nitrogen pentoxide, N 2 5 . 

Some of the more important points in connection with 
these oxides will be studied. 

Nitrous Oxide. — Nitrous oxide is a colorless gas with a 
sweet, pleasant taste and odor. "When breathed in large 
quantities it renders the patient insensible to pain, and 
if long continued will cause death. Dentists use it as an 
anaesthetic. Nitrous oxide is nearly as good a supporter 
of combustion as oxygen itself, and nearly everything 
will burn in it, uniting with the oxygen and liberating 
the nitrogen. It is made by heating ammonium nitrate 
(the ammonium salt of nitric acid). The nitrate decom- 
poses into water and nitrous oxide, as represented in the 
equation, NH 4 N0 3 = 2H 2 + N 2 0. It should be collected 
over hot water, since it dissolves to a considerable extent 
in cold water. 

Nitric Oxide. — Nitric oxide contains twice as much 
oxygen in proportion to the nitrogen as is found in ni- 
trous oxide. It is not easily decomposed, and hence is 
not a good supporter of combustion like nitrous oxide, 
although it contains so much oxygen. On the other hand, 
it combines directly with oxygen on contact with it, and 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 121 

forms a heavy red brown gas called nitrogen peroxide. 
When nitric acid is allowed to act upon copper and other 
metals part of the acid is decomposed while oxidizing the 
metal, and nitric oxide is one of the products formed. 
Nitric oxide is a colorless gas, but on coming in contact 
with the air it takes up oxygen, and nitrogen peroxide is 
the result. Thus, NO + = N0 2 . 

To generate nitric oxide the apparatus should be ar- 
ranged as for hydrogen. Scraps of copper and a little 
water are put into the flask and nitric acid added as 
needed. The nitric oxide is nearly insoluble in water and 
so may be collected over water. During the reaction the 
flask will first fill with the red nitrogen peroxide on ac- 
count of the air which is present, but this dissolves in 
the water, while the nitric oxide passes on through the 
water and up into the jars. 

In proportion to the nitrogen present there is four 
times as much oxygen in the peroxide as in nitrous oxide. 
When nitrogen peroxide dissolves in water the resulting 
solution is found to contain both nitrous and nitric acids 
in nearly equal quantity. 

The other two oxides of nitrogen, nitrogen trioxide 
and nitrogen pentoxide, contain respectively three and 
five times as much oxygen in proportion to the nitrogen 
as is found in nitrous oxide. Nitrogen trioxide is some- 
times called nitrous anhydride, since with water it forms 
nitrous acid. (An anhydride is an oxide which with 
water will form an acid.) This oxide is a blue green 
liquid. Nitrogen pentoxide is a white solid, and is the 
anhydride of nitric acid. 

Of the five oxides of nitrogen the last two are of less 
interest to the young student, perhaps, than the other 
three. Nitrous oxide, as just mentioned, is used by den- 
tists; nitric oxide and the peroxide are largely used in 
the manufacture of sulphuric acid. (See p. 107.) 



122 



A PRACTICAL CHEMISTRY 



Numerous other compounds of nitrogen might be men- 
tioned, but few of them have any place in an elementary 
study. Their connection, however, with the class of sub- 
stances known as explosives should be considered. 

Explosives. — An explosive is any substance which can 
be suddenly converted into a large volume of expanding 
gases. The sudden bursting out of these gases is an 




:^mwmim 



*mm*&~3 s 



Fig. 50. — A Blast. (By courtesy of E. I. du Pont de Nemours Powder Company.) 



explosion. Nearly all explosives contain compounds of 
nitrogen, since many of these compounds decompose sud- 
denly into large volumes of gases at high temperatures. 
The following are a few of the more common explosives: 
Black Gunpowder. — The oldest explosive in practical 
use is gunpowder. Its composition as made at various 
times and in various places has varied somewhat, but 75 
per cent potassium nitrate, 15 per cent charcoal and 10 
per cent sulphur constitute a good powder. The pow- 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 123 

dered ingredients are mixed and wet with water and 
ground together to a fine paste, which is compressed 
into cakes and dried and then broken into grains. These 
grains are glazed by rubbing with graphite and sorted 
as to size by sifting. 

When black gunpowder is fired a volume of carbon 
monoxide and free nitrogen many hundred times the vol- 
ume of the powder is formed, and the sudden expansion 
of these gases produces the force of the explosion. The 
potassium of the potassium nitrate forms with the sul- 
phur such compounds as potassium sulphate and potas- 
sium sulphide, and these together with unburnt powder 
constitute the smoke. 

Smokeless Powder. — The smoke of black gunpowder is 
so objectionable that smokeless powders are largely tak- 
ing its place. A number of different kinds of these pow- 
ders are in use, but in many respects they are much 
alike, having nitrocelluloses as their basis. 

Cellulose is a compound of carbon, hydrogen, and oxy- 
gen of very complicated structure found in the cell walls 
of all plants. Cotton is almost pure cellulose. When 
cellulose is treated with a mixture of concentrated sul- 
phuric and nitric acids, nitric acid reacts with the cellu- 
lose in several different proportions, forming several dif- 
ferent nitrocelluloses, according to the length of time, 
temperature, etc., employed. The most explosive of these 
is called guncotton. A solution of non-explosive nitro- 
celluloses in alcohol and ether is the ordinary collodion 
used in photography and as "new skin" for wounds. 
The smokeless powders are composed either of nitrocel- 
luloses alone or mixed with nitroglycerine and other ni- 
trates. In the explosion of guncotton the carbon and 
hydrogen are completely burned, and the nitrogen liber- 
ated, so that the products of the explosion are gases. 

Nitroglycerine and Dynamite. — A most violent explo- 



124 A PRACTICAL CHEMISTRY 

sive is made by pouring ordinary glycerine into a mixture 
of nitric and sulphuric acids cooled to the proper tern- 
perature. Nitroglycerine is a liquid which may be 
heated and sometimes burned without exploding, but 
which will explode when jarred by a slight shock. 

To render nitroglycerine less explosive it is absorbed 
in some porous material as sawdust, sand or infusorial 
earth, forming dynamite. By increasing the proportion 
of inactive material the resulting dynamite is made less 
explosive, and will withstand shipping and rough hand- 
ling. It may, however, be made almost as explosive as 
nitroglycerine by increasing the proportion of this in- 
gredient. 

Nitroglycerine is a single chemical compound and is 
known chemically as glycerine trinitrate, C 3 H 5 (N0 3 ) 3 . 

Picric Acid. — Picric acid is another explosive composed 
of a single chemical compound called trinitrophenol, 
C 6 H 2 (N0 2 ) 3 OH. It is made by the action of nitric acid 
on carbolic acid. It is a yellow crystalline solid which 
is also used as a dye. It is sometimes used in fireworks. 

Numerous other explosive substances exist. Among 
these may be mentioned fulminate of mercury, a dread- 
fully explosive substance used for making detonators 
(the caps by which other explosives are fired). This ex- 
plosive is very powerful and dangerous, since it is ex- 
ploded by slight shocks. 

"While the descriptions of the manufacture of explo- 
sives seem simple, the processes are very complicated, 
and require much care and many precautions; so they 
should be performed only by experienced people. 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 125 

SUMMARY 

The nitrogen of the air is inactive and combines directly with 
but few elements, but when combined with other elements 
it becomes active and useful. Nearly all compounds of 
nitrogen can be made from nitric acid and ammonia. 

Nitric Acid. — Nitric acid is usually made by the action of sul- 
phuric acid on nitrates. It is sometimes made by dissolving 
certain oxides of nitrogen in water. These oxides are formed 
by burning nitrogen in the electric arc. In nature this is 
done by lightning. In the soil nitrates are formed by the 
action of bacteria upon decaying animal and vegetable mat- 
ter in the presence of lime or alkalies. Nitric acid is a strong 
acid and a strong oxidizing agent. Nitrates are soluble in 
water. 

Ammonia. — Ammonia is a gas produced by the decomposition 
of animal and vegetable matter, and by the destructive dis- 
tillation of coal. It is very soluble in water and with water 
forms ammonium hydroxide. Ammonium compounds are 
those compounds containing the radical NH 4 . A radical is 
a group of elements forming part of a molecule. Liquid 
ammonia is used in the manufacture of ice. 

Nitrous Acid. — Nitrous acid contains less oxygen than is found 
in nitric acid. It is less stable and not so important as nitric 
acid. Its salts, the nitrites, are made by the reduction of 
nitrates. The acid may be made from its salts by decom- 
posing them with sulphuric acid. 

Oxides of Nitrogen. — Five oxides of nitrogen are mentioned in 
this chapter, each of which has characteristic properties. 
They may all be made from nitric acid or its salts. Three 
of them dissolve in water forming acids. 

Explosives. — An explosive is a substance which can be suddenly 
converted into a large volume of expanding gases. An 
explosion is the sudden bursting out of expanding gas or 
gases. Nearly all explosives contain nitrogen compounds. 
Black gunpowder is composed of potassium nitrate, char- 
coal, and sulphur. The gases formed by its explosion are 
carbon monoxide and nitrogen. 



126 A PRACTICAL CHEMISTRY 

Nitroglycerine is glycerine trinitrate, and is made by pour- 
ing glycerine into a mixture of nitric and sulphuric acids. 

Dynamite is nitroglycerine absorbed by some porous ma- 
terial to render it less explosive. 

Trinitrophenol is made by the action of nitric acid upon 
carbolic acid. 

Smokeless powders are usually made of nitrocellulose 
alone, or of nitrocellulose mixed with nitroglycerine and 
other nitrates. Numerous other explosives are known. 



REVIEW QUESTIONS 

1. What advantages are derived from the inactivity of the 
nitrogen of the air? 

2. Why do explosives usually contain nitrogen compounds? 

3. Explain how nitrates are formed in nature. 

4. Why do not nitrates accumulate more rapidly in the soil ? 

5. Why does the nitric acid in the laboratory bottles become 
colored? 

6. How may spontaneous combustion be produced by nitric 
acid? 

7. What is the action of nitric acid on various metals? 

8. What is the composition of ammonia? 

9. How do you explain the existence of ammonia in the air? 

10. Give the various names of a solution of ammonia in 
water? 

11. How does the ammonium radical differ from a metal? 

12. How is the ammonia of commerce obtained? 

13. Could any other hydroxide be substituted for calcium 
hydroxide in the laboratory method for generating ammonia? 

14. Explain the use of liquid ammonia in making ice? 

15. Why is brine used in cold storage houses? 

16. Could you use ammonium nitrite for making nitrous 
oxide? 

17. How is sodium nitrite made? 

18. How do nitrites in water indicate that the water is unfit 
for drinking? 



COMPOUNDS OF NITROGEN AND EXPLOSIVES 127 

19. Why is it not possible to discover the odor of nitric oxide? 

20. How is nitric oxide made in the laboratory? 

21. What acids can be made from nitrogen peroxide? 

22. What property of black gunpowder is objectionable? 

23. What advantage would be gained by using black gun- 
powder as a porous material in making dynamite? 

24. How does a steam boiler explosion differ from an explo- 
sion of gunpowder? 

25. Mention two different methods for causing explosives to 
explode. 



CHAPTER XIV 
FERTILIZERS 



Man is dependent upon the soil for his own food and 
for that of his animals. Year after year the plants take 
from the soil those substances necessary for their growth. 
Year after year the rains wash away more of these same 
materials than is removed by the plants. Slowly the dis- 



mm iL ','^m^k ^fm, • 



Fig. 51. — An Unfertilized Field. Without fertilizer the crop will evidently be poor. 

integration of the rocks and the decay of vegetable mat- 
ter produce more soil, but this slow process cannot fully 
replenish the loss unassisted by man. 

The substances which man must thus add to the soil 
to increase its productiveness are called fertilizers. 

Composition. — It is evident that the fertilizers must 
furnish all the mineral matter needed by the plant not al- 

128 



FERTILIZERS 



129 



ready present in sufficient quantity in the soil. The fertil- 
izer should also supply plenty of vegetable matter, which 
by decaying furnishes humus, or vegetable mold. In addi- 
tion to the carbon dioxide which plants take from the 
air and the moisture from the ground, they require in 
considerable quantity compounds containing phosphorus, 
potassium, nitrogen and calcium, and in smaller quantity 
several other elements, such as magnesium and iron, etc. 
A complete fertilizer would contain all of these, but 




Fig. 52. — A Field Fertilized with Potash. Fertilized with potash the prospects 
for a good yield are much improved. 



some soils are not in need of them all. Thus, one soil 
may be rich in phosphate, but lacking in potassium com- 
pounds, while another may need only phosphates. All 
fertilizers should be soluble in cold water, so that plants 
may be able to take them up. The making of a fertile 
soil, however, often involves more than the mere addition 
of plant food, since much can be done to improve the 
physical condition of the soil. Hard, lumpy soils may be 
rendered more porous and friable, while those which are 
too porous may have this fault corrected. Figure 54 is a 
picture of a fertile field in the time of harvest. The soil 
of this field was naturally light, sandy and porous, but 



130 



A PRACTICAL CHEMISTRY 



by the use of proper fertilizers and the plowing under of 
suitable plants it has been greatly improved. 

Sources of Fertilizers. — 1. Stable Manure. — The require- 
ments of a complete fertilizer are more nearly met in 
stable manure than in any other substance, since it is of 
vegetable origin and contains all the elements used by 
plants. 

2. Bones of Animals.- — Bones of animals contain cal- 
cium, phosphorus, and nitrogen, and when ground to a 




Fig. 53. — More Complete Fertilization of the Same Field. Only with com- 
plete fertilization can the best results be obtained. 

fine powder may be used as fertilizer. When thus used, 
however, the powdered bone must decay in the ground 
before most of its compounds become soluble and avail- 
able as plant food. This process requires much time ; 
hence it is customary to treat bones with, sulphuric acid 
before using them for fertilizer. By this means the cal- 
cium phosphate which they contain is decomposed, form- 
ing acid calcium phosphate and calcium sulphate, which 
are soluble. This mixture of calcium sulphate and acid 
calcium phosphate is sometimes called superphosphate of 
lime. The term acid calcium phosphate means that only 



FERTILIZERS 131 

part of the hydrogen of the phosphoric acid has been re- 
placed by calcium. Thus the formula of phosphoric acid 
is H 3 P0 4 and of calcium phosphate is Ca 3 (P0 4 ) 2 , while 
the formula of acid calcium phosphate is Ca(H 2 P0 4 ) 2 . 
The reaction between calcium phosphate and sulphuric 
acid may be expressed thus, 

Ca 3 (P0 4 ) 2 + 2H 2 S0 4 == 2CaS0 4 + Ca(H 2 P0 4 ) 2 
Calcium Superphosphate of lime 

phosphate 

3. Phosphate Rock. — This is an important source of 
phosphates, and is found in large quantities in South 
Carolina, Florida and Tennessee. To decompose the cal- 




Fig. 54. — A Fertile Field. A field of alfalfa in Delaware at harvest. The soil of 
this field is naturally poor but has been greatly improved by proper fertilizing 
and cultivation. 

cium phosphate which it contains it is treated with sul- 
phuric acid, as described above. 

4. Potassium Compounds. — The potassium compounds 
originally in the soil were derived from minerals contain- 
ing potassium. One of the most common of these min- 
erals is feldspar, which is a constituent of granite. The 
sources in the United States of potassium compounds 
are wood ashes, refuse of plants, refuse of sugar manufac- 
ture, wool washings and stable manure. Excepting the 



132 A PRACTICAL CHEMISTRY 

last named, the most important of these is wood ashes. 
The value of this substance as fertilizer depends not only 
upon the fact that it contains potassium, calcium and cal- 
cium phosphate, but also upon the fact that the potassium 
and part of the calcium are present as carbonates which 
serve to neutralize the acid in the soil. The importance 
of this will be discussed later. 

The greatest source of potassium is found in the min- 
eral deposits at Stassfurt, Germany. Among the min- 
erals found here are carnallit, kainit, and sylvinit, which 
are composed of chlorides and sulphates of potassium and 
magnesium. Much kainit is crushed and ground and 
used in this form as fertilizer. Its composition is quite 
variable. It contains potassium chloride, magnesium sul- 
phate, and water of crystallization, mixed with common 
salt and other substances. The percentage of potassium 
in kainit is quite low. 

Comparatively pure potassium sulphate and chloride 
can be obtained from the potassium minerals by repeated 
recrystallizations, and these purer salts are coming into 
use as fertilizers. 

In the fertilizer industries potassium salts are called 
"potash salts," and the quantity of potassium present 
in a given fertilizer is calculated as though it were potas- 
sium oxide, K 2 0, which is spoken of as "pure potash.' ' 
Thus a given sample of kainit might be said to contain 
12 per cent of pure potash. 

Many soils which would otherwise be fertile are poor 
in potassium compounds. The addition of a few hundred 
pounds of wood ashes per acre or an equivalent amount 
of potash salts to such soils renders them quite produc- 
tive. 

5. Nitrogen. — It is evident that the sources of potas- 
sium compounds and phosphates are largely mineral. 
These supplies exist in considerable quantity. The case 



FERTILIZERS 133 

of nitrogen is a little different. This element constitutes 
a large part of the air, as we have already learned. But 
in the air the nitrogen is in the free, or uncombined, con- 
dition, unsuitable for the nourishment of plants. As a 
fertilizer nitrogen should be in compounds. The nitrogen 




Fig. 55. — Nodules on the Roots of Clover. The nodules which contain the 
bacteria which change the nitrogen of the air into plant food may be seen by 
anyone who will carefully pull up a bunch of clover and wash off the soil from 
its roots. 

in the great storehouse of the air, in order to become 
available for plants, must be converted into compounds. 
In nature this is mostly accomplished through the growth 
of bacteria, though a small amount of nitrogen is caused 
to combine with oxygen by the lightning. In past ages 
the work of these bacteria has accumulated a supply of 
nitrates, especially in dry countries, which has been the 
world's store for all purposes. The great deposits of so- 



134 



A PRACTICAL CHEMISTRY 



dium nitrate (Chili saltpeter) in Chili are the most impor- 
tant, but these are being used at a rapid rate, and before 
many years will be exhausted. 

Much has been done in the past few years toward solv- 
ing the problem of producing nitrogen compounds for 
fertilizers. The first step in this direction was the dis- 
covery that nitrifying bacteria live in nodules on the 




Fig. 56. A Demonstration of Making Nitric Acid by the Electric Arc. The 

air is drawn through the arc within the bulb. Here some of the nitrogen is burned 
to nitrogen peroxide which dissolves in the water in the wash bottle, converting it 
into nitrous and nitric acids. 

roots of leguminous plants, such plants as clover, peas, 
and beans. These bacteria convert the nitrogen of the 
air into compounds capable of being used by the legumi- 
nous plants. The nitrogen compounds thus formed are 
stored in the roots and stems of the plants, and when 
these decay in the soil furnish nitrogen fertilizer for 
otfter crops, such as wheat and corn. Thus the farmer, 
if he wishes to increase the nitrogen content of his soil 
before planting wheat, may plant the ground in peas in 



FERTILIZERS 



135 



the early summer and later plow these under to furnish 
fertilizer for the wheat which he plants in the fall. The 
decay which the nitrogen-containing plants undergo in- 
volves the growth of several kinds of nitrifying bacteria, 
which produce ammonia, nitrous acid and finally nitric 
acid. Another class of bacteria, known as denitrifying 
bacteria, liberate nitrogen from nitrogen compounds and 
permit it to escape into the air. Explosions also liberate 




Fig. 57- — Diagram Showing Transformations of Nitrogen in Nature. 



nitrogen. This cycle of nitrogen in nature is represented 
diagrammatically in Figure 57. In Norway nitrogen is 
being burned in large electric arcs and the oxides thus 
formed are converted into nitric acid and calcium nitrate 
Another method now employed in Europe makes use of 
the fact that nitrogen will unite with highly heated cal- 
cium carbide (a compound which we have studied in 
connection with acetylene (see p. 77), forming a sub- 
stance known as calcium cyanamide. This substance is 
a valuable nitrogen fertilizer. 

6. Lime. — The term lime when used in connection 



136 A PRACTICAL CHEMISTRY 

with fertilizers has several meanings. Usually it means 
calcium oxide, CaO, or calcium hydroxide, Ca(OH) 2 , or 
calcium carbonate, CaC0 3 , but sometimes it is used in a 
more general way to represent calcium in any of its com- 
pounds. 

The benefits derived from the use of lime (whether in 
the form of oxide, hydroxide or carbonate) are of three 
kinds, physical, chemical and biological. 

The physical changes which may be produced in soils 
have already been mentioned (see p. 129). Both the ce- 
menting of sandy soils, making them less porous and 
more capable of holding water, and the rendering of clay 
soils more porous and friable, can be accomplished by 
the addition of lime. 

Chemically, lime acts in many ways upon the soil: (a) 
It corrects the acidity by neutralizing the acids present, 
(b) It converts phosphates of iron and aluminium into 
calcium phosphate, and by its action upon certain sili- 
cates renders potassium present in the soil available for 
plant food, (c) Lime hastens the decomposition of vege- 
table matter and at the same time neutralizes the acids 
formed by its decomposition. When considerable quan- 
tity of organic matter is present in the soil, as for exam- 
ple when stable manure and lime are used together, or 
when peas have been plowed under, this hastening of the 
decomposition of the vegetable matter is an advantage, 
but an excessive use of lime destroys too rapidly the 
humus of the soil, (d) Lime also decomposes mineral 
salts which have a toxic effect upon the growth of plants. 

From the biological side the most marked benefit from 
lime is found in the fact that it makes possible the 
growth of nitrifying bacteria, since these organisms can- 
not live in an acid soil. Hence, such plants as clover, 
which depend upon the nitrifying bacteria living in 
nodules upon their roots for their supply of nitrogen, 



FERTILIZERS 137 

grow well when fertilized with lime. Lime is also used 
to prevent certain plant diseases, but often with doubtful 
success. 

Only a very small part of the lime used in fertilizers 
serves as plant food, but it will be seen from the fore- 
going that as a "soil amendment" it is very important. 

Since our very existence depends upon the fertility of 
the soil, too much can hardly be said as to the importance 
of this subject of fertilizers. Future generations will be 
compelled not only to be less wasteful of the old forms 
of fertilizers, such as stable manure, slaughter house 
refuse, sewage, etc., but they will feel the necessity of 
devising new methods of fertilizing, striving for the time 
"when every rood of land maintains a man." 



SUMMARY 

Fertilizers. — Fertilizers are those substances which must be added 
to the soil to increase its productiveness. They include both 
plant food and soil amendments. 

Compounds of phosphorus, nitrogen, potassium, calcium, 
magnesium and other elements are included in plant food. 

The making of a fertile soil involves the improvement of 
its physical condition as well as the addition of plant food. 

Sources of Fertilizers. 

1. Stable manure, which includes most of the elements of a 

perfect fertilizer. 

C (a) Bones f Rendered available 

2. Phosphates from -j (b) Phosphate i by treatment with 

L (c) Rocks [ sulphuric acid 



3. Potassium Compounds: 
In the United States from 



(a) Wood ashes 

(b) Refuse of plants 

(c) Refuse of sugar manufacture 
I (d) Wool washings 



138 



A PRACTICAL CHEMISTRY 



In Germany from 



4. Nitrogen: 

(a) Nitrates in the earth 
from 



(a) Carnallit 

(b) Kainit 

(c) Sylvinit and other minerals 

1. Electrical discharges 

2, Action of bacteria in decay of 
nitrogenous matter 



(b) Growth of leguminous plants 



(c) Manufactured prod- 
ucts 



1. Nitric acid from air 

2. Calcium nitrate from air 

3. Calcium cyanamide 



5. Lime as 



(a) Calcium Car- 
bonate 

(b) Calcium Ox- 
ide 

(c) Calcium hy- 
droxide 



" 1. Correct acidity of soil 

2. Convert other phosphates 
into calcium phosphate 

3. Hasten the decomposition 
of organic matter 

Used 4. Decompose injurious min- 
' to eral salts 

5. Promote the growth of 
nitrifying bacteria 

6. Prevent diseases of plants 

7. Improve the physical 
condition of soil 



REVIEW QUESTIONS 



1. Why are not all soils equally fertile? 

2. How can you improve a clay soil? 

3. How improve a sandy soil? 

4. When a fertilizer is required, why is not stable manure 
always used? 

5. Is superphosphate of lime made from bones a better fer- 
tilizer than superphosphate of lime made from phosphate rock? 

6. Enumerate the sources of potassium compounds. 



FERTILIZERS 139 

7. Explain how the use of wood ashes may help to increase 
the nitrogen content of the soil. 

8. Why would not ground feldspar make a good fertilizer? 

9. What is meant by the expression, "Kainit contains 12 
per cent of pure potash"? 

10. Devise an experiment by which you could tell the needs 
of a given soil? 

11. Name three kinds of plant food found in the atmosphere. 

12. Why are nitrogenous fertilizers expensive? 

13. How do thunder storms make the soil more fertile? 

14. Explain the origin of the nitrates in the soil. 

15. Tell the various ways in which you might add nitrogen 
fertilizer to a wheat field. 

16. Tell one way in which the electric furnace may be used 
in connection with the manufacture of a fertilizer. 

17. How do explosions increase the cost of fertilizers? 

18. -What treatment would you give a light, sandy soil to 
improve its physical condition? 

19. How does lime liberate plant food in the soil? 

20. Suggest some ways of conserving the fertilizer resources 
of our country. 



CHAPTER XV 
FOOD 

At least three times per day the normal- body feels the 
need of food. Probably no one substance contains every- 
thing necessary for the complete nourishment of the adult 
human being, but milk comes nearer to being a perfect 
food than any other material. 

A perfect food should contain fats, nitrogenous mat- 
ter, carbohydrates, and all the mineral compounds nor- 
mally found in the body. 

Minerals. — The mineral matter required in our food in- 
cludes quite a number of elements, most of which are 
in comparatively small quantities and are taken with 
other foods almost without thought on our part. Thus 
from meat, milk, vegetables, and grains we obtain phos- 
phates, sulphates, chlorides, fluorides, and other salts of 
such metals as calcium, sodium, potassium, iron, mag- 
nesium, and lithium. Only in the case of common salt, 
sodium chloride, and water do we take mineral foods 
directly and as such. 

Carbohydrates. — To this class belong starch and sugars, 
as well as many substances which are not foods. The 
name, carbohydrates, was derived in accord with the sup- 
position that all chemical compounds of this kind are 
composed of carbon with hydrogen and oxygen in the 
proportion to form water. 

Starch is found in quantity in all kinds of grain and 
potatoes and to some extent in almost all plants. It is 

140 




I 




142 A PRACTICAL CHEMISTRY 

obtained mostly in this country from corn. The process 
is almost entirely mechanical. The corn is steeped for 
some hours in dilute sulphurous acid, and then broken 
up in water. The portion of the corn grain known as 
the germ, being lighter than the other portions, floats on 
the water and is removed by skimming. The remainder 
of the grain is ground up with water and the starch is 
washed out. It is then separated from the larger par- 
ticles by means of sieves. The water with the finely 
divided starch floating in it flows slowly over long tables, 
where the starch settles out. Later it is collected and 
dried. Starch is composed of grains, which under the 
microscope show definite markings. Starch grains de- 
rived from different sources, as from corn and potato, 
are of different sizes and appearance and may be easily 
recognized under the microscope. 

Under ordinary conditions starch is not soluble in 
water, but when a mixture of starch and water is heated 
the grains swell and burst and a paste is formed. Starch 
reacts with iodine, forming a blue compound. This may 
be made a test for the presence of starch. 

"When starch is mixed with very dilute mineral acids 
and boiled for some hours it is converted into glucose. 
The syrup thus formed is filtered through animal char- 
coal to remove coloring matter, the acid is neutralized 
with sodium carbonate, which with the hydrochloric acid 
forms common salt, and the liquid concentrated by evap- 
oration in vacuum kettles. These vacuum kettles are 
large, closed-up boilers from which the air and steam are 
being constantly pumped. By means of these the syrup 
is evaporated more rapidly and at a lower temperature, 
thus removing the danger of scorching the sugar. The 
corn syrups so widely sold at present for table use are 
composed largely of glucose syrup. 

Besides glucose another class of compounds known as 




Fig. 59. — A Diagram Showing Section of a Vacuum Pan. 



144 A PRACTICAL CHEMISTRY 

dextrines can be made from starch. These compounds 
are formed in the digestion of starch and are later 
changed into sugars. The dextrines mixed with water 
form a sticky paste which is much used as mucilage. 
Dextrine for this purpose is made either by heating starch 
with a very little nitric acid to 110° C, or without the 
acid to a somewhat higher temperature. It is evident 
that dextrine is often formed in the cooking of starchy 
foods ; for example, in the Crust of bread. 

The Sugars. — Closely related to the starches are the 
sugars, which constitute a large group of chemical com- 
pounds. 

A number of the sugars occur in foods, some of them 
in sufficiently large quantity to justify their mention here. 

The ordinary sugar usually used in cooking is known 
in chemistry as cane sugar, or sucrose, whether obtained 
from the sugar cane, sugar beet or sugar maple. It is 
found also in many fruits, nuts and flowers. Its formula 
is C 12 H 22 1:l . 

The manufacture of ordinary granulated cane sugar is 
a large industry and includes many processes. The juice 
is obtained by crushing the cane between rolls. It is next 
bleached with sulphur dioxide ; after this it is boiled with 
lime. The lime forms insoluble compounds with many 
of the impurities, which either sink to the bottom or may 
be skimmed from the top. The liquid, after filtering, is 
ready to be evaporated by passing through a series of 
vacuum pans. A very thick syrup is thus formed which 
deposits crystals on cooling. The crystals are separated 
from the syrup by centrifugal machines — revolving cylin- 
ders with sieve walls. The liquid is thrown out of the 
cylinders as they revolve, while the crystals are retained 
by the meshes. The crystals are then dried by currents 
of warm air which are passed through the granulators in 
which the drying is done. Much sugar is not refined 



FOOD 145 

when made, but is shipped as crude sugar to the sugar 
refineries of our large cities. Here it is dissolved in water 
and lime and phosphoric acid added. This mixture forms 
insoluble calcium phosphate, which settles to the bottom 
of the containing tank, and carries down with it much of 
the impurities in the sugar solution. The solution is then 
filtered, first through canvas bags and then through bone- 
black. The remaining treatment is the same as that de- 
scribed above for making granulated sugar directly from 
sugar cane. 

The remaining liquid, after the sugar crystals have 
been removed, is molasses. 

Much sugar is also made from sugar beets. In this 
process the beets, after being thoroughly cleaned, are 
sliced up. The slices are then dropped into tanks of 
water, where the sugar dissolves, forming what is known 
as "diffusion juice." This liquid is purified, first by fil- 
tering and then by a treatment with lime and carbon 
dioxide. The carbon dioxide precipitates the lime as 
calcium carbonate, which carries down with it many of 
the impurities. The juice, after further filtering, is 
bleached with sulphur dioxide. The remaining steps of 
the process are similar to those in the making of sugar 
from sugar cane. 

Maple sugar and maple syrup are made by boiling 
down the sap of the sugar maple, which is obtained by 
boring small holes into the trees and driving in "spiles." 
These are small wooden or metal troughs through which 
the sap runs into buckets placed below. The sap runs 
most freely in the spring, when the nights are cold and 
the days are mild. The maple syrup requires little or no 
purifying, apart from the removal of mineral matter 
which settles to the bottom. The peculiarly pleasant taste 
of maple sugar is due to small quantities of organic com- 
pounds which it contains. Cane sugar, when heated care- 



146 A PRACTICAL CHEMISTRY 

fully, will melt. at 160° C. to a clear liquid, but when 
heated higher begins to turn brown, owing to the forma- 
tion of caramel, a dark, almost tasteless, substance. Car- 
amel candy, no doubt, contains some of this compound. 
Boiled with dilute acids, a very different change occurs; 
one molecule of cane sugar takes up a molecule of water 
and separates at the same time into two molecules of two 
other sugars known as glucose and fructose, as indicated 
by the equation: 

Ci 2 H 22 O l:L + H 2 = C 6 H 12 6 + C G H 12 6 

These sugars are not so sweet as cane sugar, and on 
this account, when sour fruit is cooked for some time 
with sugar, the sugar loses some of its sweetening power. 

Glucose is also called dextrose and grape sugar. It is 
found in many fruits mixed with fructose, which is also 
called fruit sugar and levulose. Both glucose and fruc- 
tose have a composition represented by the formula 
CeH-^Og, but the atoms are not arranged in the same way 
in the two molecules. As previously stated (see p. 142), 
much glucose is made from starch. It is used largely in 
the manufacture of candy and syrups. 

Another sugar of some importance is lactose, or milk 
sugar, which is obtained by evaporating the whey, or 
watery portion of milk. It is used in the manufacture 
of infant foods and in medicines. 

Fermentation. — If the juice of fruit or other liquids 
containing glucose are exposed to the air at moderate 
temperature, changes of a peculiar sort soon take place. 
The sugar present is decomposed, and alcohol and carbon 
dioxide are formed. The change is indirectly brought 
about by a microorganism called yeast. This process is 
alcoholic fermentation, and yeast is called a ferment. 
The yeast produces a chemical compound called zymase, 
which belongs to the class of compounds known as en- 



FOOD 147 

zymes. The zymase, acting as a catalytic agent, produces 
the fermentation. 

Many other substances besides glucose undergo fermen- 
tation, and yeast is but one of a number of ferments. 
The ferments are microscopic plants. Two other very im- 
portant fermentations besides the alcoholic are the acetic 
and lactic acid fermentations. These will be discussed 
later. 

Alcoholic fermentation is conducted on a large scale 
for the purpose of obtaining alcohol and alcoholic bev- 
erages. 

Ordinary, or ethyl, alcohol (C 2 H 5 OH) is but one of a 
large number of alcohols known to chemists. Small quan- 
tities of many of these are formed along with ordinary 
alcohol during alcoholic fermentation of glucose. Al- 
coholic fermentation continues until the solution contains 
about 14 per cent alcohol. Alcohol of this concentration 
kills the ferment. The ethyl alcohol may then be sep- 
arated from most of the water and other alcohols by frac- 
tional distillation. The mixture of other alcohols thus 
separated is known as fusel oil. 

Under proper conditions starch also can be made to 
undergo this fermentation, but a concentrated solution 
of cane sugar will kill the ferment. Advantage is taken 
of this fact in preserving fruit, since a thick syrup of 
cane sugar is used to protect the fruit from fermentation. 
Canned fruits, however, are further protected by being 
kept in airtight cans. In this way the yeast and other 
ferments floating in the air are prevented from entering 
the fruit, while those already present are killed by cook- 
ing the fruit. 

Alcoholic beverages are of two classes, those which are 
merely fermented and those which after fermentation are 
distilled. The latter contain a larger percentage of al- 
cohol. 



148 A PRACTICAL CHEMISTRY - 

Beer and wines are fermented liquors. In making beer, 
barley is first converted into malt by letting it sprout 
and then killing it by heating. The malt is then fer- 
mented with hops. Beer contains from 3 to 7 per cent, 
of alcohol. Wines run as high as 20 per cent. 

Whisky, brandy, and rum are examples of distilled 
liquors. Whisky is made by fermenting corn, wheat, or 
barley, and distilling off the alcohol -from the mash thus 
formed. The whisky is next "rectified" to remove the 
other alcohols, which are usually spoken of as fusel oil. 
Coloring matter is then added. Whisky contains about 
50 per cent, of alcohol. 

Ethereal Salts, or Esters. — W x e have learned in a pre- 
vious chapter (see p. 36) that bases will neutralize acids 
and thus form salts. Alcohols, in some respects, resemble 
the bases which we have previously studied. While they 
will not directly neutralize acids as sodium hydroxide 
does, they can be made to react with acids and form neu- 
tral compounds called ethereal salts, or esters. These 
are unlike ordinary salts in that they are not crystallized 
compounds, but are often liquids or semisolids, like lard 
and tallow. Oil of wintergreen, or methyl salicylate, is a 
good example of the liquid ethereal salts. Many of the 
flavors of fruits and odors of flowers are due to the pres- 
ence of these ethereal salts. Artificial perfumes and 
flavorings are made of ethereal salts. For example, the 
synthetical banana oil is amyl acetate, that is, it is made 
from amyl alcohol and acetic acid. Since there are so 
many alcohols and acids, the number of etheral salts is 
almost unlimited. 

Fats and Oils. — Glycerine is a substance which most 
people would not think of as an alcohol, and yet chemi- 
cally it belongs to that class. Its formula is C 3 H 5 (OH) 3 . 
It combines with many acids, particularly those carbon 
acids which contain a large number of carbon atoms, 



FOOD 149 

forming ethereal salts of a peculiar sort. These are fats 
and oils. The great variety of these natural fats and oils 
is due, not only to the large number of acids with which 
glycerine naturally combines, but also to the fact that 
various mixtures of the various simple fats often occur 
in foods. 

The acids from which fats are most often derived are 
oleic, palmitic, and stearic, and the fats formed from 
these acids with glycerine are known respectively as 
glyceryl oleate, or olein; glyceryl palmitate, or palmitin, 
and glyceryl stearate, or stearin. The formulae of these 
compounds are quite complicated ; thus we have for stearin 
C 3 H 5 (C 17 H 35 C0 2 ) 3 , in which the C 3 H 5 is the glyceryl radi- 
cal derived from glycerine and the three C 17 H 35 C0 2 are 
from three molecules of stearic acid, C 17 H 35 C0 2 H. Some 
examples of fats are lard, tallow, butter, olive oil, and 
palm oil. Tallow contains much stearin, while in palm 
oil and olive oil there is olein. Butter is a good example 
of a fat containing a large number of ethereal salts ; 
among these are glyceryl salts of palmitic, oleic, stearic, 
butyric, caproic, and capric acids. 

The fats used for food are derived not only from meat 
but also from grains, nuts, seeds, et cetera. Thus from 
the germs removed from corn, as described in starch-mak 
ing, a corn oil is expressed. Likewise we have cottonseed 
oil, olive oil, et cetera. 

Nitrogenous Food. — Food of this sort is usually classed 
under the head of proteids, which is a general term used 
to represent a great variety of different substances, such 
as the albumen of eggs and the gluten of wheat, the 
casein of milk, gelatin, and numerous other substances. 
The term protein is now also given to these compounds, 
though this word formerly had a different meaning. Pro- 
teids are of the utmost importance in food. The body 
cannot exist without them, though one may live for some 



150 A PRACTICAL CHEMISTRY 

time without other kinds of food. They are among the 
most complicated substances known to chemistry, their 
molecules containing many hundreds of atoms. They 
contain carbon, hydrogen, oxygen, and nitrogen, a small 
percentage of sulphur, and sometimes phosphorus. Ow- 
ing to their complexity and the ease with which they de- 
compose, the proteids are studied with great difficulty, 




Fig. 60. — A Modern Cow Stable. Since the importance of clean milk has become 
recognized, much care is taken at the best dairy farms to keep the cow stables 
free from filth and dust. They are also well ventilated and well lighted. 

and as yet comparatively little is known of their chem- 
istry. 

Everyone is familiar with the coagulation of the al- 
bumen of eggs during boiling ; a similar change occurs in 
many other proteids also. Such chemical substances as 
acids and alcohol often produce a like result. This is 
illustrated by the curdling of milk. 

Milk. — Milk has previously been mentioned as ap- 
proaching a perfect food. This is true, in that it contains 



FOOD 151 

fat and nitrogenous matter, a carbohydrate and mineral 
substances. The first two are held in suspension as an 
emulsion, and the others are in solution in the water or 
whey of the milk. 

Many kinds of bacteria find in milk the conditions un- 
der which they can grow and multiply with great rapid- 
ity. On account of this and the great carelessness and 
lack of cleanliness employed in handling milk, it often 
happens that many thousands of bacteria are found in 
each cubic centimeter of milk. The most common of 
these bacteria is the lactic acid ferment, which acts upon 
the carbohydrate or lactose in the milk, converting it into 
lactic acid, and the milk becomes "sour." This does not 
occur suddenly. The bacteria fall into the milk from the 
air, the sides of the cow, and from the clothing of the 
milker, and begin at once to bring about the fermentation 
of the lactose, but some time elapses before sufficient acid 
is formed to give the milk a sour taste. The action of the 
other kinds of bacteria in milk often produces many other 
changes which may escape the notice of the ordinary ob- 
server, but sometimes the odors and flavors produced are 
noticeable to all. 

The action of all bacteria in milk may be delayed by 
low temperature, and completely stopped by antiseptics. 
Among the preservatives sometimes used in milk may be 
mentioned hydrogen peroxide, formaldehyde, and sali- 
cylic acid. The method most used at present is pasteur- 
izing, which consists in subjecting the milk to a high tem- 
perature for a short time and then chilling it. 

When milk is carefully pasteurized and then protected 
from the air it will keep for some time, since the process 
kills nearly all the bacteria, but does not kill the spores 
After a few days these spores, which correspond to seed 
of ordinary plants, will develop into a new crop of bac- 
teria. The process of pasteurizing is named from the bar 



152 



A PRACTICAL CHEMISTRY 



teriologist Pasteur, who discovered that it is possible to 
kill most of the disease-producing bacteria without injur- 
ing the milk. The temperature and length of time em- 
ployed by different milk dealers vary somewhat; some 
hold the milk at a temperature of 165° F., or 65° C, for 




Fig. 61. — The Pasteurizing Room of a Large City Milk Dealer. The milk is 
heated in large boilers like the one at the right of the picture. It is then cooled 
in water coolers, one of which is shown at the front of the picture. 



twenty minutes ; others use a lower temperature, but con- 
tinue the heating for a longer time. 

When milk is bought unpasteurized the consumer may 
easily pasteurize it at home by carefully heating it in a 
bottle suspended in a vessel of water, or in a "double 
boiler." Care is required that the temperature does not 
go above 165° F. "When milk is heated sufficiently high 
so that all bacterial life is destroyed it is said to be 
sterilized. This can be done by keeping the milk at 
212° F. for two hours. This process gives the milk an 



FOOD 



153 



unpleasant cooked taste and produces undesirable changes 
in its constituents. It is much better, however, to avoid 
the bacteria in the first place by careful handling than 
to use any means to remove them after they have been 
admitted to the milk through lack of cleanliness. Dairy- 




Fig. 62. — Many Thousands of Clean Bottles Ready to Receive the Pasteur- 
ized Milk. 



men and milk dealers are giving much thought to this 
subject. (Figures 60 and 62.) 

The nitrogenous matter in milk is casein. It is held in 
suspension in the whey of the milk, but may be coagulated 
by acids, alcohol, or rennet. This coagulation of the 
casein or the formation of curds is most often seen in 
sour milk, in which it has been brought about by the 
lactic acid. The thick milk thus formed is called clabber. 
If clabber is scalded with boiling water, the casein is 
toughened and may be removed from the whey by filter- 
ing through a cloth. The curd thus obtained is some- 



154 A PRACTICAL CHEMISTRY 

times used as "pot cheese," or "cottage cheese." "When 
dried, casein finds use as a substitute for gum in applying 
color to paper and other similar uses. It is also combined 
with lime to form a white solid for the manufacture of 
such small articles as buttons. Casein as it exists in milk 
will form insoluble compounds with several of the min- 
eral poisons, particularly lead acetate (or sugar of lead), 
mercuric chloride (or corrosive sublimate), and silver ni- 
trate, hence the use of large draughts of milk in cases of 
poisoning with these substances. 

In cheese-making the casein is coagulated with rennet, 
which is a substance extracted from the linings of calves' 
stomachs. 

Our best American Cheddar cheese is made from whole 
milk. The milk is brought to a temperature of about 
30° C, and a sufficient quantity of rennet added to coagu- 
late the casein in a few minutes. The curds are then cut 
with knives and kept warm for some time till the whey 
separates and the casein begins to toughen. The curd is 
next put into cheese cloth in a filter press the shape and 
size desired for the cheese and the whey squeezed out. 
After removal from the press the cheese is stored in a 
drying house for some months while the ripening takes 
place. This is one of the most important processes in the 
manufacture of cheese and is brought about by bacteria 
of various kinds. If improper ferments enter the cheese, 
decomposition sets in and the product is spoiled. On the 
other hand, the Various desirable flavors depend as much 
upon the introduction of the proper bacteria as upon the 
material from which the cheese is made. It follows, then, 
that the various kinds of cheeses derive their peculiar 
qualities from the differences in materials and bacteria 
and methods of manufacture used. "Whole milk cheeses 
contain much of the fat of the milk. Filled cheeses are 
made from milk from which much of the butter fat has 



SANITARY FAUCET 

EXTRA HEAVY TINWARE 

REVERSIBLE FLOAT 



CENTER BALANCED BOWL 



ONE PIECE OETACHED SPINOLE 



HELICAL TOOTH SPUR PINION 
AND WORM WHEEL GEARS 



BRONZE REVERSIBLE WORM WHEEL 
FRAME JOINING SCREW 
OPEN SANITARY BASE 




Fig. 63. — The De Laval Milk Separator. The milk is poured into the supply can 
at the top and passes down through the faucet into the bowl which is revolving at 
a high speed. Here the heavier portions of the milk move to the outer part of the 
bowl and by the aid of the discs ascend to the top and. pass out into the skim- 
milk tube. The cream, which Is lighter than the skim-milk, moves toward the 
middle of the bowl and passes out at a hole at the "cream screw" and then 
flows away through the other tube into a cream can (not shown in the cut). 



156 A PRACTICAL CHEMISTRY 

been removed and other fats added in its place. In some 
kinds of cheese other kinds of milk, as goats' milk, are 
employed. 

Cream consists of whey, some casein, and the fats of 
milk. It is separated from milk either by permitting the 
milk to stand for some hours in a cool place, when the 
cream, being lighter than the milk, comes to the top and 
is skimmed off, or by centrifugal force in a machine 
called a separator. This machine consists of a rapidly 
revolving bowl through which the milk flows. The milk 
takes up the rotary motion of the bowl. The heavier por- 
tions of the milk are whirled to the wall of the bowl and 
escape through one pipe, while the lighter cream remains 
at the center of the bowl and passes out by a different 
exit. 

If the cream is to be made into butter, it must first pass 
through a ripening process, which is accomplished by 
bacteria. In the best creameries the cream is first steril- 
ized and then a pure culture of the proper sort of bacteria 
added. As in the case of cheese, the bacteria produce 
the flavor. The cream is churned, after the ripening is 
complete, by agitating it in a churn. This causes the 
little particles of fat to separate from the casein and 
come together in masses. After this the butter is thor- 
oughly washed, and salted if desired. 

Butter, as previously stated,, is a mixture of fats. It 
contains also some casein, water, and small quantities of 
the various substances formed by bacterial action. But- 
ter becomes "strong" on account of the decomposition 
of the fats, particularly glyceryl butyrate, with the libera- 
tion of free acids. 

Vinegar. — Many substances enter into foods which ap- 
parently have little food value. Among these may be 
mentioned vinegar, which is dilute, impure acetic acid. 
The acetic acid fermentation has already been mentioned. 




Fig. 64. — Butter Making. The old way. 




Fig. 65.— Butter Making. The new way. 



158 A PRACTICAL CHEMISTRY 

It is brought about by a ferment known commonly as 
"mother of vinegar," a microscopic plant which may be 
seen in great masses in old vinegar casks, and resembles 
dirty brown rags. This ferment in the presence of oxy- 
gen of the air oxidizes alcohol into acetic acid thus : 

C 2 H 5 OH + 2 = CH 3 C0 2 H + H 2 
(Alcohol) (Acetic 

acid) 

The best vinegar is supposed to be that made from cider. 
Apples are crushed to a pulp and the juice squeezed out 
by pressure, forming cider. Alcoholic fermentation at 
once begins (the yeast plants being present on the apples) 
and the sugar of the cider is converted into alcohol and 
carbon dioxide. This fermentation continues for some 
time, increasing the amount of alcohol in the cider. Soon 
the acetic acid fermentation sets in, and the alcohol is 
converted into acetic acid as just described. The product 
is vinegar. 

Vinegar is sometimes made from low-grade wines and 
other alcoholic liquors by subjecting them to the acetic 
acid fermentation. This is often accomplished by letting 
the alcoholic liquid drip through beech shavings which 
have been previously mixed with mother of vinegar. Air 
is at the same time allowed to come in contact with the 
alcohol, and the fermentation goes on rapidly. 

The vinegars thus far described are colored. Colorless 
vinegar can be made by diluting pure acetic acid. 

Acetic Acid. — Acetic acid, CH 3 C0 2 H, is made in large 
quantities in the destructive distillation of wood. The 
wood is heated in closed iron retorts and acetic acid is 
one of the products distilled off. The acid is neutralized 
with lime, forming calcium acetate, and this is purified 
by crystallization and evaporation of volatile impurities. 
The calcium acetate is next mixed with sulphuric acid, 







Fig. 66. — Making Cider. The apples are ground in a mill on the floor above and then 
passed through a chute into canvas or burlap bags as shown on the car at the right 
of the picture. The car with its load is then pushed into the hydraulic press, which 
squeezes the cider from the apple pulp. 

















.;.; 










j *'• 










\m 


■ * 




:' f i \ \' .;.■'} 




i ■ - : 














■ 1 








^"-jjj ' 










) ' 












P^j22yui 


1 ~^J 






Ill 


Kppp^JH J, 




: - 


j 


1 Mi 














WKBfSi 







Fig. 67. — Tanks in Which the Cider Changes to Vinegar. 



160 A PRACTICAL CHEMISTRY 

and the acetic acid is distilled off while the~ calcium sul- 
phate thus formed remains behind in the retort. 

Acetic acid is a liquid with strong acid odor. It freezes 
at 17° C. when free from water, and on this account the 
water-free acid is known as glacial acetic acid. 

In addition to its use in vinegar, acetic acid is de- 
manded for many other things, such as the manufacture 
of white lead, dyestuffs, acetates, and medicines. 

Bread. — While it is far from possible for us to discuss 
all the various substances used in food, bread is such a 
common article of diet that it seems to require special 
mention. It is also appropriate to bring it up at this 
time on account of its connection with fermentation. 

As a food wheat bread stands high on account of the 
number of constituents which it contains. It is rich in 
starch, contains nitrogenous matter in the form of gluten, 
and is fairly well supplied with mineral matter, particu- 
larly phosphates. In addition to this bread usually car- 
ries considerable fats. 

So far as methods of manufacture are concerned, dif- 
ferent kinds of bread may be divided into three classes : 
(1) those that are raised, that is, made light and puffy 
by means of ferments; (2) those in which the raising is 
done by means of chemical mixtures called baking pow- 
ders; (3) those which are not raised. Bread which is not 
raised is called unleavened bread. 

In ordinary wheat bread the flour, salt, lard, sugar, and 
water are mixed with yeast, and permitted to stand until 
the alcoholic fermentation producing alcohol and carbon 
dioxide has furnished enough gas to lighten the dough. 
After being molded into shape the bread is baked. Dur- 
ing the baking the alcohol is driven out by the heat, while 
the carbon dioxide by its expansion further raises the 
bread until it begins to crust over and harden. If the 
bread is not baked promptly after it is raised, the second 



FOOD 161 

or acetic acid fermentation will begin, and vinegar will 
be formed in the bread from the alcohol. The heat will 
not drive out the acetic acid and the bread will be sour. 

When baking powders are used for raising bread, cakes, 
and biscuit, the puffing up is again accomplished through 
the expansion of carbon dioxide, but the dioxide is ob- 
tained from the decomposition of a carbonate and not 
from alcoholic fermentation. 

Carbonates are the salts of carbonic acid, H 2 C0 3 , and 
the particular carbonate used in baking powders is acid 
sodium carbonate or sodium bicarbonate, which is com- 
monly called baking soda. The formula of acid sodium 
carbonate is NaHC0 3 . With acids, carbonates decompose 
with the formation of carbon dioxide, a new salt of the 
base which was in the carbonate and water. A study of 
the reaction between baking soda and hydrochloric acid 
will make this clear. 

HNaC0 3 + HC1 = NaCl + H 2 + C0 2 
(Acid sodium (Hydrochloric (Sodium (Water) (Carbon 
carbonate) acid) chloride) dioxide) 

Baking Powders. — Baking powders are mixtures of acid 
sodium carbonate or baking soda and some solid acid sub- 
stance, with a little corn starch to keep the mixture dry. 
When used in dough, the water present brings the acid 
substances and the baking soda together, and a reaction 
similar to the one described above takes place. The car- 
bon dioxide formed is caught by the dough, which is 
baked at once, and the heat expanding the gas causes it 
to raise the bread or cake. 

Baking powders generally contain baking soda, but 
several different acid substances are used. Among these 
may be mentioned cream of tartar or monopotassium tar- 
trate, tartaric acid, acid calcium phosphate, and alum. 

It is necessary that we study some of these various 



162 A PRACTICAL CHEMISTRY 

compounds used in baking powders. The starting point 
in the manufacture of baking soda is common salt or 
sodium chloride, and one of the methods much employed 
is called the Solvay process. A cold, concentrated brine is 
made of salt and water. Into this ammonia gas, NH 3 , is 
passed until the solution is saturated, and this is fol- 
lowed with carbon dioxide, which is also bubbled through 
the solution until it will dissolve no more. A reaction 
takes place which may be represented by the following 
equation : 

NaCl + NH 3 + C0 2 + H 2 = NaHC0 3 + NH 4 C1 
(Sodium (Ammo- (Car- (Water) (Acid sodi- (Ammo- 
chloride) nia) bon um car- nium 
dioxide bonate) chloride) 

Acid sodium carbonate is almost insoluble in ammonium 
chloride solution, and hence is precipitated as a solid 
powder. It is filtered from the solution and dried. 

The by-product, ammonium chloride, is not lost. By the 
method described in the ammonia chapter it is treated 
with calcium hydroxide and the ammonia recovered to 
use again. The reaction involved may be expressed thus : 

2NH 4 C1 + Ca(OH) 2 = CaCL, + 2NH 3 + 2H 2 
(Calcium (Calcium 
hydroxide) chloride) 

The lime for this reaction is obtained by burning lime- 
stone, which is calcium carbonate, and decomposes ac- 
cording to the equation : 

CaC0 3 = CaO + C0 2 

(Calcium (Calcium (Carbon 

carbonate) oxide) dioxide) 

The C0 2 may be used in the Solvay process. It is thus 
seen that the only materials required for the process are 
salt, limestone, water, and coal. 



FOOD 163 

Tartaric Acid.— Tartaric acid is an acid found in many- 
kinds of fruits. It is a solid crystallized substance which 
may be manufactured synthetically or gotten from its 
salts. One of the most common of the latter is argol, or 
tartar, which collects on the inside of wine casks. "When 
this is purified it is called cream of tartar, or acid potas- 
sium tartrate, or monopotassium tartrate. This means 
that the tartaric acid has been but half neutralized by 
potassium. The formulae of tartaric acid and cream of 
tartar are C 2 H 2 (OHC0 2 ) 2 H 2 and C 2 H 2 (OHC0 2 ) 2 HK, from 
which it may be seen that in cream of tartar one potas- 
sium atom has replaced one hydrogen atom of the acid, 
leaving one hydrogen atom yet to be replaced, so that 
the resulting compound still has acid properties. The 
acid calcium phosphate used in baking powders is of a 
similar nature; here only part of the hydrogen of the 
phosphoric acid has been replaced by calcium, and the 
resulting substance has acid properties. 

Alum. — Alum, which is often used in place of acids in 
baking powder, however, is not, strictly speaking, an acid 
substance, but it undergoes with baking soda a compli- 
cated reaction in which both alum and soda are decom- 
posed and carbon dioxide is liberated. 
. The word alum should be plural, for there are a number 
of alums. This is perhaps the first illustration we have 
had of double salts. If two salts formed of two different 
metals with the same acid are dissolved in the same solu- 
tion and then permitted to crystallize frequently they 
will unite to form one crystal. In the case of the alums 
the salts are those of sulphuric acid, the metals or bases 
are usually sodium, potassium, lithium, or ammonium in 
the one sulphate, and aluminium or chromium in the 
other. The metal or base in one sulphate has a valence 
of one, as sodium, potassium, etc., and in the other has a 
valence of three, as aluminium and chromium. The crys- 



164 



A PRACTICAL CHEMISTRY 



tals are regular octahedral — eight-sided— solids with 
twelve molecules of water of crystallization. Thus, if a 
mixture of sodium sulphate and aluminium sulphate solu- 
tions is allowed to crystallize, sodium aluminium alum is 
formed, while if the solution contained potassium sulphate 
and chromium sulphate the resulting crystals would be 
potassium chromium alum, which is usually called chrome 
alum. The formula of chrome alum may be written 
KCr(S0 4 ) 2 -f 12H 2 0, and that of ordinary soda alum 
NaAl(S0 4 ) 2 + 12H 2 0. 

It will be noticed that in the formula the water of 
crystallization is separated from the rest of the molecule 
by a plus sign; it is also customary to use a period for 
the same purpose. The reaction between ordinary alum 
and acid sodium carbonate may be expressed : 



3NaHC0 3 

(Sodium 

acid 

carbonate) 



+ AlNa(S0 4 ) 2 
(Alum) 



+ 2Na 2 S0 4 
(Sodium 
sulphate) 



+ 



Al(OH), 

(Aluminium 
hydroxide) 

3C0 2 
(Carbon 
dioxide) 



Beverages. — We have previously made mention of some 
of the alcoholic liquors. These practically have no food 
value and cannot be classed as foods. In addition to 
these are several other beverages often taken with food, 
though some of them are of questionable value. Among 
these are tea, coffee, chocolate, and cocoa. Mention must 
also be made of mineral waters. 

Coffee and Tea. — Coffee and tea are beverages too fa- 
miliar to require much introduction. We find in both of 
them the same alkaloid, caffeine, which is a compound 
chemically related to uric acid. Caffeine is of no food 



FOOD 



165 



value, but has a stimulating effect upon portions of the 
brain. In large quantities it is doubtless injurious not 
only to the nervous system, but also to the heart and 
digestive organs. Besides this alkaloid a small amount 
of food is found in coffee, but the quantity in tea is too 
small to be considered. By alkaloids we mean organic 




Fig. 68. — Leaves, Flowers and Pods of the Cocoa Plant. Illustration copy- 
righted and loaned by the Walter Baker Co., Ltd. 



bases. They are compounds of carbon, hydrogen, nitro- 
gen, and sometimes oxygen, which react with acids as 
ammonia does, forming salts. They often have a bitter 
taste. 

Chocolate and Cocoa. — Chocolate and cocoa are both 
derived from the seed of various species of trees of the 
Theobroma genus. In composition the two substances 
differ primarily in the percentage of fats, proteids, and 




Fig. 69. — Cocoa Pods Containing the Cocoa Beans. The pods are from nine 
inches to a foot long. The cocoa beans are about the size of almonds and are em- 
bedded in tissue like the seeds in a watermelon. The beans are removed from the 
pods and are subjected to a sort of fermentation, after which they are dried in 
the air. Courtesy of Rockwood & Co. 









FOOD 



16- 



carbohydrates which they contain. If a considerable por- 
tion of the fat is removed from chocolate, the remaining 
portion is practically cocoa. Both substances are of the 
highest food value, though chocolate is more difficult for 
some people to digest on account of the large amount of 
fat which it contains. 







* 








\ 


* iBI 














j9 




























ii 


'-S? 4 


ST 


s-i 


1 



Fig. 70. — Moulding Cocoa. When the cocoa beans are received at the factory they 
are first freed from foreign matter (sticks, stones, etc.). They are then roasted 
and the hulls removed. After this they are crushed and the germ, or part of the 
bean which produces the sprout, is removed. The beans are next ground to a thin 
paste. The chocolate thus formed is subjected to pressure by means of a hydraulic 
press and part of the fat removed. This fat is cocoa butter. After the removal 
of part of the fat, the remaining product is cocoa. Illustration copyrighted and 
loaned by the Walter Baker Co., Ltd. 



A chemical substance called theobromine is found in 
the seeds from which chocolate and cocoa are obtained. 
It is chemically related to caffeine and uric acid, but 
does not seem to have any injurious effect upon the heart 
and other organs. 

Mineral Waters. — Under this heading should be in- 



168 A PRACTICAL CHEMISTRY 

eluded not only the waters of mineral springs, such as 
Baden-Baden, Saratoga, and the White Sulphur Springs, 
but also artificial waters. 

In the United States there are more than eight thou- 
sand mineral springs, the waters from several hundreds 
of which are bottled and sold. These various waters con- 
tain a great variety of mineral ingredients, among which 
may be mentioned sodium carbonate and sulphate, mag- 
nesium, calcium and lithium carbonates, sodium bromide 
and iodide, silica, hydrogen sulphide, carbon dioxide, and 
many other substances. 

The manufactured mineral waters are supposed to be 
of the same composition as those of the noted springs 
from which they are named. For example, the manu- 
factured vichy is intended to have the composition of the 
water from Vichy. These waters are usually charged 
with carbon dioxide and are said to be carbonated. 

Soda Water. — Soda water is a simple form of manufac- 
tured mineral water made by forcing carbon dioxide into 
cold water under pressure. The carbon dioxide unites 
with water, forming carbonic acid. An acid containing 
an excess of its anhydride in solution is called a fuming 
acid. Chemically, then, soda water may be said to be 
fuming carbonic acid. Numerous flavors are mixed with 
this plain soda. 

SUMMARY 

Perfect Food. — A perfect food should contain fats, nitrogenous 
matter, carbohydrates, and all the mineral compounds nor- 
mally found in the body. 

Minerals. — The minerals are found in other foods and in water, 
and include phosphates, sulphates, chlorides, fluorides, and 
other salts of calcium, sodium, potassium, iron, manganese, 
lithium, et cetera. 

Carbohydrates. — Carbohydrates include starches, dextrines, and 



FOOD 



169 



sugars. Starch is made from potatoes, corn, and other grains 
by processes which are largely mechanical. Starch may be 
converted into dextrine and glucose by boiling with acids. 
It may also be made into dextrine by heat alone. 



Cane 

Beet 

Found ini Nuts 

Flowers 
_ Fruits 
Decomposed by heat, 
caramel being one 
of the products 
Boiled 



f Sucrose or 
cane sugar 



There are many 
sugars. Those 
mentioned are : 



C12H22O11' 



with 
acids 
forms 



Glucose 
Fructose 



CeH^Oe 



Lactose or 
milk sugar 



Glucose or dex- 
trose or grape 
sugar 



Fructose or 
levulose or 
fruit sugar 



Found in milk. Fer- 
ments to lactic acid 

Used in medicine and 
in infant foods 

Found in grapes and 

other fruits 
Made from starch 

by boiling with 

dilute acid 
Ferments forming al- 
„ cohol 

Found in fruit 
Ferments forming 
alcohol 



Fermentation. — Fermentation is a chemical change indirectly 
brought about by some living microorganism; the micro- 
organism usually produces some chemical compound which, 
acting as a catalytic agent, causes the chemical change. The 
ferments mentioned in this chapter are: 



170 A PRACTICAL CHEMISTRY 



f Glucose 
Fermenting < Fructose 



Alcoholic (Yeast) 



Starch, etc. 



f Ethyl alcohol 



- Forming -l Carbon dioxide 
[ Other alcohols 

Acetic acid I Fermenting ethyl alcohol 
I. Forming acetic acid 

Lactic acid j Fermenting lactose 
I J ormmg lactic acid 

Alcoholic beverages are of two classes: those which are 
merely fermented, as beer and wines, and those which after 
fermentation are distilled, as whisky and brandy. 

Ethereal Salts. — Ethereal salts are salts made from alcohols in- 
stead of metallic hydroxides. They are often liquids or semi- 
solids. They include artificial perfumes and flavorings and 
the fats and oils. 

Fats are glyceryl ethereal salts of the fatty acids. The 
glyceryl radical, C 3 H 5 , is derived from the alcohol glycerine, 
C 3 H 5 (OH) 3 . Such fats as tallow are mixtures of the sim- 
ple fats, stearin, palmitin, and olein. Fats used as food are 
derived from meats, nuts, and grains. 

Nitrogenous Foods. — Nitrogenous foods include proteids or pro- 
tein — such substances as the albumen of eggs and the gluten 
of wheat. The body cannot exist without such compounds. 
They are complicated and studied with difficulty. 

Milk. — Milk contains fats, nitrogenous matter, lactose, mineral 
substances, and water. Bacteria grow rapidly in milk and 
convert the lactose into lactic acid, making the milk sour 
and coagulating the casein. The bacteria also produce many 
other undesirable changes. The growth of bacteria is hin- 
dered by low temperature, by preservatives, and by pasteur- 
izing. 



FOOD 



171 



Cheese. — Cheese is made by coagulating the milk with rennet, 
pressing out the whey, or watery part of the milk, and al- 
lowing the curd to dry and ripen. 

Cream. — Cream contains the fats of milk, together with whey 
and some casein. It is removed from milk by means of a 
machine called a separator. Butter is made by churning the 
cream, which causes the particles of fat to collect in lumps. 
Butter is a mixture of glyceryl ethereal salts of fatty acids. 
It also contains a little casein and water. 

Vinegar. — Vinegar is dilute acetic acid produced by the acetic 
acid fermentation of alcohol in cider and wines. It is also 
made by diluting pure acetic acid obtained from the destruc- 
tive distillation of wood. 



Bread. — There are three 
classes of bread 



r Those kinds raised by ferments 
producing alcohol and carbon 
dioxide 
Those kinds raised by baking pow- 
ders producing carbon dioxide 
Those not raised 



Baking Powders Baking 

powders are made of 



Acid sodium 
carbonate 

Starch 

An acid sub- 
stance 



Cream of tartar 
Tartaric acid 
Acid calcium- 
phosphate 
Alum 



Acid sodium carbonate, also called sodium bicarbonate 
and baking soda, is made by the Solvay process. Starch is 
used to keep the mixtures dry. Alums are double sulphates 
of aluminium and an alkali metal or ammonium. The 
alums all crystallize with twelve molecules of water of crystal- 
lization. When alums react with acid sodium carbonate 
aluminium hydroxide and carbon dioxide are formed instead 
of aluminium carbonate; hence it is possible to use alums 
for acid substances in baking powders. 



172 



A PRACTICAL CHEMISTRY 



Beverages T h e 

beverages men- 
tioned in this 
chapter include 



Alcoholic 
drinks 



r Beer 
Wines 
Whisky 
Brandy, etc. 



Not food but 
injurious 



{Containing 
caffeine and 
a little food 



Sometimes 
injurious to 
nerves and 
digestive or- 
gans 

Chocolate J Containing . much food and 
and cocoa I not usually injurious 
Mineral wa- f Containing mineral food, and 
ters I of medical value 



REVIEW QUESTIONS 



1. What should a perfect food contain? 

2. How do the mineral foods differ from other foods? 

3. What is the name carbohydrate intended to signify? 

4. How would you get starch from potatoes? 

5. How could you tell cornstarch from potato starch? 

6. How is the presence of starch detected? 

7. What becomes of the hydrochloric acid used in making 
glucose? 

8. How is it possible to boil water at a temperature below 
100° C? 

9. How is it possible to heat water to 110° C? 

10. What is starch paste? 

11. What substance is found in the crust of bread? 

12. Why is the sugar made from sugar beets called cane 
sugar? 

13. Why do we highly value crude maple sugar and yet re- 
quire that sugar made from cane and beets be refined? 

14. How can you manufacture mucilage from starch? 

15. How could granulated sugar be made from maple sugar % 

16. How does molasses differ from corn syrup? 






FOOD 173 

17. When cane sugar is boiled with cream of tartar what 
reaction occurs? 

18. How can lactose be obtained as a by-product in cheese- 
making? 

19. What change occurs in solutions containing glucose when 
exposed to the air? 

20. How does zymase differ from the other catalytic agents 
thus far studied? 

21. What is meant by fusel oil? 

22. What are hops? 

23. Is malt ever used as food? 

24. Why does cider never contain more than 14 per cent 
alcohol? 

25. How do ethereal salts differ from ordinary salts? 

26. Name an ethereal salt often found in candy, and state its 
composition. 

27. How does gtycerine chemically resemble ethyl alcohol? 

28. Why does it require three molecules of stearic acid to 
react with one molecule of glycerine? 

29. How do you account for the great number and variety of 
the fats? 

30. What kind of food is of most importance? 

31. Why would you expect the proteids to decay before the fats ? 

32. Mention some common illustrations of the coagulation of 
proteids. 

33. How does an emulsion differ from a solution? 

34. How does milk sometimes spread disease? 

35. What effect does the souring of milk have upon the quan- 
tity of sugar which it contains? 

36. Why does sour milk become clabber? 

37. How does pasteurized milk differ from "raw" milk? 

38. What is the best way to keep milk? 

39. Why should not cows be fed while being milked? 

40. What constituents of milk do not remain in the cheese? 

41. Why do not all samples of milk give the same percentage 
of cream? 

42. Why are various samples of butter so different in prop- 
erties? 



174 A PRACTICAL CHEMISTRY 

43. How do the fermentation processes in bread-making re- 
semble those in making vinegar? 

44. Write two chemical equations involved in the making of 
glacial acetic acid. 

45. Why is glacial acetic acid so named? 

46. Explain how the various acid substances react with the 
acid sodium carbonate in baking powders. 

47. What by-product of the Solvay process cannot be used 
again in the process? 

48. Tell how each of the other by-products is saved and used 
again. 

49. Why does it require more cream of tartar than of tartaric 
acid to react with a given quantity of sodium acid carbonate? 

50. Write the names and formulae of four alums. 

51. Why do alums change to a white powder when heated ? 

52. Why are coffee and tea not good for children? 

53. Which is of more food value, chocolate or cocoa? 

54. Does soda water have any food value? 

55. What food can the body obtain from mineral waters? 






CHAPTER XVI 
SOME FAMILIAR COMPOUNDS USED IN MEDICINE 

We do not expect in this chapter to make a study of 
the thousands of substances used in medicine — mineral 
compounds, drugs derived from plants, secretions of ani- 
mal glands, and synthetic compounds — but rather to 
familiarize the young chemist with a few compounds in 
everyday use. We shall take these up under several 
general headings as anaesthetics, et cetera, and try also to 
learn a little of some of the elements which they contain. 

Antiseptics. — The science of medicine has undergone 
many changes in past years, but probably nothing has 
wrought greater change than the discovery of bacteria 
and of their connection with disease. Along with this 
came the discovery of those substances which we call 
antiseptics and disinfectants. Among these we will con- 
sider hydrogen peroxide, formaldehyde, carbolic acid, 
mercuric chloride, iodoform, sulphur dioxide, boric acid, 
and borax. 

Hydrogen Peroxide. — We have learned before of this 
oxide and of its relation to water. We will recall the 
formula of water, H 2 0, and that of hydrogen peroxide, 
H 2 2 , and notice that the apparent difference in the mole- 
cules is an atom of oxygen. Hydrogen peroxide is a thick 
syrup-like liquid, which decomposes when pure with 
great violence. Pure hydrogen peroxide, however, is 
almost unknown, and that which is ordinarily sold is 
nearly 97 per cent, water. It is usually made by the 

175 



176 A PRACTICAL CHEMISTRY 

action of phosphoric acid, H 3 P0 4 , or of dilute sulphuric 
acid, H 2 S0 4 , in the presence of phosphoric acid upon 
barium peroxide, Ba0 2 . The hydrogen of the sulphuric 
acid and barium of the barium peroxide exchange places, 
as may be shown by the following equation: 

H 2 S0 4 + Ba0 2 = H 2 2 + BaS0 4 . 

(Barium sulphate) 

Dilute hydrogen peroxide decomposes slowly at ordi- 
nary temperature, each molecule forming one molecule 
of water and one atom of oxygen in the nascent condition : 

H 2 2 = H 2 + 0. 

In the presence of any substance capable of being easily 
oxidized this reaction is quite rapid. All are familiar 
with the frothing produced when hydrogen peroxide is 
poured upon a sore. Hydrogen peroxide is sometimes sold 
under the name of dioxogen. 

Formaldehyde. — A very different though equally active 
antiseptic is formaldehyde. Its formula is HCHO. It is 
the simplest of a number of aldehydes, which are com- 
pounds formed by the oxidation of the alcohols. Formal- 
dehyde is made by the oxidation of methyl, or wood alco- 
hol, CH3OH. A 40 per cent, solution of formaldehyde gas 
in water is used as a disinfectant of rooms where there 
has been a contagious disease. This is conveniently done 
by pouring the solution on quicklime. The reaction be- 
tween the lime and the water produces sufficient heat to 
vaporize the formaldehyde, which penetrates to every part 
of the room and kills the bacteria. The chemical action 
of this disinfectant is that of a reducing agent, that is, it 
abstracts oxygen from other things and is oxidized to 
formic acid. 

Carbolic Acid. — For general purposes of disinfecting 
probably nothing is better than carbolic acid, since it can 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 177 

be used in such a variety of ways. The quantity required, 
however, is greater than of most antiseptics. 

The term carbolic acid is usually applied to a water 
solution of phenol, though sometimes phenol is called car- 
bolic acid. Phenol, C 6 H 5 OH, is a white crystalline sub- 
stance with a characteristic odor. It is not a strong acid, 
but will destroy the skin. It is a violent poison. 

Mercuric Chloride. — In mercuric chloride, HgCl 2 , or 
corrosive sublimate, or bichloride of mercury, as it is 
often called, we have still another type of antiseptic, in 
that it is strictly a mineral compound. 

It is much used in surgery, and while not very soluble 
in cold water it is so active that a comparatively dilute 
solution serves as an excellent antiseptic wash. Like all 
other soluble mercury compounds, corrosive sublimate is 
a dreadful poison for which a good antidote is raw white 
of egg. This should be followed by an emetic. 

In connection with this compound it is right that we 
should take a little notice of the element, mercury, the 
only metal we have liquid at ordinary temperature. Mer- 
cury occurs in nature combined with sulphur in the min- 
eral, cinnabar, which is mercuric sulphid, HgS. From 
this ore mercury is easily obtained by roasting with air, 
by which the sulphur is oxidized and passes on 2 as the 
gas, sulphur dioxide, while the mercury vapor is cooled 
and condensed. Mercury finds many uses in connection 
with science, while many of its compounds are employed 
in medicine. The use of the fulminate in explosives has 
been mentioned. 

Mercury has two valences ; the higher or -ic valence is 
two, as may be seen from the formula of mercuric 
chloride. Mercuric oxide, HgO, is the red compound in 
which Priestley discovered oxygen. 

Iodoform. — Iodoform, CHI 3 , is the yellow powder 
which has been used so much in dressing wounds, and the 



178 A PRACTICAL CHEMISTRY 

disagreeable odor of which is so familiar to all. This 
compound may be made by warming a mixture of dilute 
alcohol, sodium carbonate, and iodine. In fact, this reac- 
tion is used as a test for alcohol. Sodium carbonate is 
dissolved in the solution to be tested, and iodine (a little 
at a time) is added and the mixture is warmed. The 
iodine disappears, and, on cooling the solution, a yellow 
crystalline precipitate of iodoform is formed. The chem- 
ical changes involved are too complicated to discuss here, 
but the resulting compound is of interest to us in another 
way. It may be recalled that under hydrocarbons we 
studied methane, or marsh gas, which has the . formula 
CIT 4 . Now if we write I 3 in place of three hydrogen 
atoms of CH 4 , we have CHI 3 , the formula of iodoform. 
Moreover, it is possible to substitute in methane three 
iodine atoms for three of its hydrogen ; hence iodoform is 
called in chemistry tri-iodomethane. 

Sulphur Dioxide. — When either sulphur or a sulphide 
is burned or roasted in the air, sulphur dioxide, S0 2 , is 
formed. This is a suffocating gas which has been much 
used to disinfect rooms after cases of contagious diseases, 
but is not so effective as formaldehyde. It probably has 
no effect when perfectly dry, but with moisture it forms 
sulphurous acid, H 2 S0 3 , as shown in the equation, 

S0 2 + H 2 = H 2 SO, 
Sulphur dioxide. Sulphurous acid. 

Sulphurous acid is a strong reducing agent. Its action 
as a disinfectant may be partly due to this, and partly to 
its uniting with other substances. This acid is also mu^h 
used as a preservative in canned fruits and vegetables. 
Its use as a bleaching agent will be discussed later. 

Boric Acid. — This compound, which is also called 
boracic acid, has the formula, H 3 B0 3 . It is found in con- 
siderable quantity in Tuscany, where it issues as vapor 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 179 

with steam from cracks in the earth. Large tanks have 
been built around these steam jets and filled with cold 
water, in which the steam and acid condense. The solu- 
tion thus formed is concentrated by evaporation, and the 
boric acid is deposited as white crystals. In several parts 
of the world mineral borates occur. This is particularly 
true in Death Valley, California, where a calcium borate 
occurs. From this borate both boric acid and borax are 
manufactured. The calcium borate, which is called cole- 
manite, has the formula CagBgOn + 5H 2 0. This is mixed 
with water and sodium carbonate in the right proportion 
in large tanks and heated for some hours. A reaction 
takes place, in which solid calcium carbonate is formed 
and borax passes into solution. The calcium carbonate is 
removed by means of a filter-press, and the borax crystal- 
lizes out of the solution on cooling. 

To make boric acid, sulphuric acid is added to the 
borax solution, and a reaction represented by the follow- 
ing equation takes place : 

Na 2 B 4 7 + H 2 S0 4 +■ 5H,0 = Na 2 S0 4 + 4H 3 B0 3 
(Borax) (Sulphuric (Sodium (Boric 

acid) Sulphate) acid) 

On cooling, the boric acid crystallizes out. 

Boric acid is a mild, non-poisonous disinfectant, partic- 
ularly valuable in lotions. Borax, the chemical name of 
which is sodium tetraborate, Na 2 B 4 7 + 10H 2 O, is of not 
much importance as an antiseptic, but is of much value 
in the manufacture of soap and for other technical pur- 
poses which we shall discuss in a subsequent chapter. 

Anaesthetics. — The important part played by antiseptics 
in the development of medical science has gone hand in 
hand with the advancement of surgery, and this in turn 
has required anaesthetics. Under the term anaesthetics 
may be included all those substances which render a part, 



180 A PRACTICAL CHEMISTRY 

or all, of the body insensible to pain. The action of these 
substances varies greatly in every particular. The effect 
of some lasts for a short time, of others for quite a long 
period. They are administered by inhalation, injection, 
and by spraying upon the part to be affected. Although 
numerous experiments had been made and much knowl- 
edge gained concerning anaesthetics and their action, 
their practical use did not begin till 1846, when Dr. W. T. 
G.' Morton, a Boston dentist, used ether. 

Chloroform. — Chloroform, CHC1 3 , is a- volatile liquid 
with a pleasant, ethereal odor. It is non-inflammable and is 
a solvent for a number of substances. When inhaled it 
soon produces insensibility, and is poisonous when taken 
into the stomach. 

By comparison of the formula of chloroform with that 
of iodoform it is seen that the two are quite similar, the 
only difference being that in one we have chlorine and 
in the other iodine. 

One method of manufacturing emphasizes this similar- 
ity of structure : alcohol and bleaching powder are heated 
together and the chloroform distills off. The bleaching 
powder furnishes the chlorine and base corresponding to 
the iodine and sodium carbonate in the making of iodo- 
form, which has been described. The purest chloroform is 
made by decomposing chloral with an alkali. Chloroform 
is one of the most important anaesthetics in use ; another 
of equal importance is ether. 

This substance is closely related chemically to alcohol, 
from which it is easily made. It is very volatile and in- 
flammable, and much care is required in handling to pre- 
vent dangerous fires. Like chloroform, it is a good solvent 
for many substances like fats and oils. It is somewhat 
soluble in water, but very much more soluble in alcohol. 

As stated above, ether is made from alcohol. At least 
two reactions are involved. The first is accomplished by 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 181 

slowly pouring a quantity of alcohol into a little more 
than twice its weight of concentrated sulphuric acid. The 
result is a sort of acid ethereal salt, known as ethyl 
sulphuric acid, thus: 

C 2 H 5 OH + H 2 S0 4 = C 2 H 5 HS0 4 + H 2 
(Ethyl (Ethyl sul- 

alcohol) phuric acid) 

It will be noticed that the ethyl radical, C 2 H 5 , of the 
ethyl alcohol takes the place of one hydrogen atom of 
the sulphuric acid. This ethyl sulphuric acid is next put 
into a distilling apparatus so arranged that after the acid 
is heated to its boiling point more alcohol may be slowly 
added as the ether distills over. The second reaction is 
represented by the equation : 

C 2 H 5 HS0 4 + C 2 H 5 OH = H 2 S0 4 + (C 2 H 5 ) 2 0. 
(Ethyl sul- (Ethyl 

phuric acid) ether) 

By adding more alcohol the process may be continued 
until a considerable amount of ether has distilled over. 

The ether which distills over is more or less mixed with 
alcohol, from which it can be freed by distillation and by 
washing with water. 

Like alcohol, ether is one of a series of compounds, and 
the ethers are named like the alcohols. "We have methyl 
alcohol, ethyl alcohol, propyl alcohol, et cetera, and 
methyl ether, ethyl ether, et cetera, and, again, just as 
the common alcohol is ethyl alcohol, so the common ether 
is ethyl ether. 

Nitrous Oxide. — In dental surgery until quite recently 
this anaesthetic has held sway. In the last few years, 
however, other things are being introduced to take the 
place of nitrous oxide. Nitrous oxide, N 2 0, was called 
laughing gas by Davy, who studied its composition and 



182 A PRACTICAL CHEMISTRY 

properties. It has a sweet odor and taste, and when in- 
haled in small quantities produces a sort of intoxication. 
Breathed in large quantities its effect is insensibility, 
which lasts for a short time. It should be avoided by peo- 
ple with weak hearts. 

Nitrous oxide is not a combustible gas, but freely fur- 
nishes oxygen for the combustion of other things. It is 
easily made by decomposing ammonium nitrate by heat. 

NH 4 N0 3 = N 2 + 2H 2 
(Ammonium (Nitrous 
nitrate) oxide) 

Like many other gases, it is much more soluble in cold 
water than in hot, and should, therefore, be collected over 
hot water. 

Local Anaesthetics. — Among the other anaesthetics in 
common use are those that are used in minor surgery and 
are applied to the part where the operation is to be per- 
formed. Some of these, by their rapid evaporation, ren- 
der the part numb with cold and hence insensible, while 
others produce a temporary paralysis. Among the latter 
is cocaine. It is obtained from the coca plant, and is a 
very dangerous and deadly alkaloid. 

A recently discovered anaesthetic is stovaine, a com- 
pound of very complicated composition. It is used by in- 
jection with a hypodermic needle into the spinal cord, 
rendering all nerves which pass off from the spinal cord 
below this point insensible to pain while the patient re- 
mains perfectly conscious. 

Alkaloids. — We have mentioned caffeine and cocaine as 
alkaloids. These are a class of chemical compounds con- 
taining nitrogen and acting as bases. They will unite 
with acid just as ammonia does to form salts. ■ Most of 
these compounds are dangerous poisons which find some 
place in medicine. The dangers in using them arise not 






SOME FAMILIAR COMPOUNDS USED IN MEDICINE 183 

only from their being poisons, but also from the fact that 
many of them produce a habit demanding their continued 
use, resulting often in some serious derangement of the 
system. Among these compounds we have, besides those 
mentioned, codeine and morphine, which are obtained 
from opium; atropine, which is sometimes used partially 
to counteract the injurious effects of morphine, when the 
drug is required; nicotine, found in tobacco ; quinine; 
stryclinine; emetine; and many others. 

We cannot emphasize too strongly the need of care in 
the use of alkaloids. Few of them should be taken except 
by direction of a physician. 

Tinctures. — Solutions formed by dissolving medicines 
in alcohol are called tinctures. Many medicines are dis- 
pensed in this form. 

A few of these tinctures are familiar to all, such as 
laudanum, which is a tincture of opium ; tincture of iron, 
made by dissolving ferric chloride in alcohol; Jamaica 
ginger ; tincture of arnica and tincture of iodine, which 
are sometimes used for sprains. 

In connection with the tincture of iodine a. few words 
concerning iodine itself are in place. 

This element belongs to the same family as chlorine, 
i. e., it is chemically related to chlorine, and has many 
properties in common with that element. It is likewise 
related to bromine and to fluorine. These four elements — 
fluorine, chlorine, bromine, and iodine — are sometimes 
called halogens, meaning "sea salt producers." Com- 
pounds of iodine are found in sea water and in Chili salt- 
peter, and from these compounds iodine is obtained. 

In very small particles, as one sees it after painting the 
skin with tincture of iodine, this element is of a brown 
color. In crystals it is from gray to black, and when 
heated and converted into a vapor it is violet. 

Iodine evaporates in the air, dissolves slightly in water, 



184 A PRACTICAL CHEMISTRY 

and freely in alcohol or in a solution of potassium iodide. 
It forms a number of acids similar to those of chlorine; 
first among these is the one corresponding to hydrochloric 
acid. Hydriodic acid, HI, however, is not so stable a com- 
pound as hydrochloric acid and is of 'comparatively little 
importance. Its salts, however, particularly potassium 
and sodium iodides, are used in medicine. 

Iodine is liberated by a reaction similar to that by 
which chlorine is obtained from its compounds ; that is, by 
the use of an oxidizing agent in the presence of sulphuric 
acid. 

2NaCl + Mn0 2 + 2H 2 S0 4 = 2C1 + MnS0 4 
(Sodium (Manganese (Manganese 

chloride) dioxide) sulphate) 

+ 2H 2 + Na 2 S0 4 . 
(Sodium (Sodium 

sulphate) sulphate) 

2NaI + Mn0 2 + 2H 2 S0 4 = 21 + MnS0 4 + 2H 2 
(Sodium 
iodide) 

+ Na 2 S0 4 . 

Bromides. — Bromides, which are also much used in 
medicine and in photography, are salts of hydrobromic 
acid, HBr, just as iodides are salts of hydriodic acid and 
chlorides of hydrochloric acid. The element bromine, 
from which hydrobromic acid is made, is a heavy, dark, 
red-brown liquid which easily evaporates to a heavy 
brownish red vapor. Bromine is dreadfully corrosive to 
the skin and mucous membranes. It reacts violently with 
many elements and with the metals forms bromides. Bro- 
mine is liberated from bromides by a reaction similar to 
the one just given for iodine. 

Emetics. — In cases of poisoning and for other purposes 
it is sometimes necessary to give medicines which will 
produce vomiting. These are called emetics. One of the 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 185 

safest is ipecac, which is derived from the root of a South 
American plant. Ipecac contains the alkaloid emetine. 

Zinc Sulphate. — Zinc sulphate, ZnS0 4 , the zinc salt of 
sulphuric acid, is sometimes used for the same purpose. 
Alum is also used, as well as several other compounds. 

Cathartics. — A very large number of substances are 
used for this purpose. A few of these are simple chemical 
compounds which should be described. 

Calomel. — This is mercurous chloride, HgCl. It is an 
insoluble white powder that forms as a precipitate when- 
ever a soluble chloride or hydrochloric acid is put into 
solution with any mercurous salt. Ammonia changes it 
to a black powder, 

Magnesium Sulphate, Milk of Magnesia, and Magnesium 
Citrate. — These are three compounds of the metal mag- 
nesium much used for this same purpose. Magnesium 
sulphate, MgS0 4 , is a crystalline powder called Epsom 
salt, from the fact that it is found in the waters of the 
noted springs at Epsom, England. 

Milk of magnesium is the white, milk-like magnesium 
hydroxide, Mg(OH) 2 . Magnesium citrate is usually sold 
in solution charged with carbon dioxide. It is a salt of 
citric acid. This acid is found in lemons, and may also 
be made synthetically. 

Caustics. — Under this heading may be mentioned, be- 
sides sodium and potassium hydroxide, lunar caustic, 
which is silver nitrate, AgN0 3 , melted and cast into sticks, 
and zinc chloride, ZnCl 2 . These are used to remove ob- 
jectionable tissues. 

Antidotes. — These are medicines or other substances 
used to counteract poisons. These substances must be of 
such a nature as to react chemically with the poison 
and destroy its nature or convert it into an insoluble sub- 
stance, or the antidote may overcome the poison by acting 
physiologically in such a way as to oppose it. As ex- 



186 A PRACTICAL CHEMISTRY 

amples of the first, alkalies and carbonates are given as 
antidotes for acids, since these will neutralize the acids, 
or, conversely, acids would be given to neutralize alkalies. 
The antidote for mercuric chloride (jaw white of eggs) 
has already been mentioned. This is an illustration of 
forming an insoluble compound with the poison. The 
albumen of the egg is coagulated by the mercuric chloride 
and holds the poison in an insoluble clot, which may then 
be removed by an emetic. Another illustration of the 
same thing is the use of freshly precipitated ferric hy- 
droxide in cases of poisoning with arsenic compounds such 
as white arsenic and Paris green. The iron forms an in- 
soluble compound with the arsenic. 

For oxalic acid chalk or any powder form of calcium 
carbonate is given. Here we have an example of both 
neutralizing the acid and forming an insoluble compound 
of calcium oxalate. 

For carbolic acid Epsom salts serves as an antidote, 
since it reacts with the acid and renders it harmless 
through the formation of magnesium sulpho-carbolate. 

Part of the poisonous effect of carbolic acid and, in fact, 
of acids and alkalies in general, is their corrosive action 
on the membranes of the throat and stomach. To over- 
come this sweet oil is often used. On the other hand, oils 
must be avoided if phosphorus from matches, rat poisons, 
and the like has been taken, since phosphorus is soluble 
in fats and oils. 

For phosphorus the first treatment should be an emetic 
or the stomach pump. Generally speaking, the stomach 
pump has the first place in treating a case of poisoning, 
or, if this is not at hand, an emetic. Often both are de- 
sirable. 

Many alkaloids, narcotics, and similar poisons are 
treated with antidotes to counteract them by their physio- 
logical action. The use of atropine with morphine is an 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 187 

illustration. Strychnine and alcohol are often used in 
this way as antidotes. 

SUMMARY 



The medical substances mentioned in this chapter may be classi- 
fied as follows: 1. Antiseptics, 2. Anaesthetics, 3. Alkaloids, 
4. Tinctures, 5. Emetics, 6. Cathartics, 7. Caustics, 8. Anti- 
dotes. 





r 


'H 2 2 






Hydrogen 


Used as a 3 per cent, solution 




peroxide « 


made from barium peroxide 
and sulphuric acid and phos- 
phoric acid 

-HCHO 

Used as a 40 per cent, solution 




Formalde- 


with quicklime to disinfect 




hyde 


rooms 
It acts as a strong reducing 
. agent. 


Antiseptics (sub- 


Carbolic acid < 


' C 6 H 5 OH, phenol 
Used as a solution 


stances which 






kill bacteria) in-^ 




It is a poison 
^HgCl 2 


elude the follow- 








Used as an antiseptic wash 


ing: 


Mercuric 


Is very poisonous 




chloride "* 




"Hg 






Contains 

- 


A liquid metal 






mercury 


Its ore is cin- 






** 


nabar 






r CHI 3 






A yellow crystalline solid 






f Alcohol, 




Iodoform - 


Made from -j iodine and 

[ a base 
Chemical name is tri-iodo- 






L methane 





188 



A PRACTICAL CHEMISTRY 







'S0 2 
Made by burning sulphur 




Sulphur 


Dissolves in water forming 


Antiseptic sub- 


dioxide 


sulphurous acid. 


stances which 




Used' also as a preservative 


kill bacteria in- - 




- and bleaching agent. 


elude the follow- 




' H 3 B0 3 


ing. 




Found in nature 




Boric acid - 


Made from borax and other 
borates 




*■ 


Non-poisonous 



Anaesthetics. — Anaesthetics are substances which render the body 
insensible to pain. The more common anaesthetics include 
chloroform, ether, nitrous oxide, cocaine, and stovaine. 

Chloroform (CHC1 3 ) is made by distilling a mixture of 
alcohol and bleaching powder, or a mixture of an alkali 
and chloral. It is a non-inflammable liquid with a pleasant 
odor. It is taken by inhalation. 

Ether (C 2 H 5 ) 2 is a volatile, inflammable liquid, made by 
distilling a mixture of alcohol and sulphuric acid. It is 
also administered by inhalation. Its chemical name is ethyl 
ether, which distinguishes it from the other members of the 
ether series. 

Nitrous oxide (N 2 0) is a gas with sweet odor, used in dental 
surgery. It is made by heating ammonium nitrate. 

Cocaine is an alkaloid used as a local anaesthetic. 

Stovaine is a recently discovered anaesthetic. 
Alkaloids. — Alkaloids are organic bases containing nitrogen. 
They are poisons which produce a habit demanding their 
continued use. Some of them are codeine, atropine, mor- 
phine, nicotine, quinine, strychnine, and emetine. 
Tinctures. — Tinctures are solutions formed by dissolving medi- 
cal substances in alcohol. In connection with the tincture 
of iodine the other halogens are mentioned. Several com- 
pounds of bromine and iodine are used in medicine. Bro- 
mine is a corrosive, red brown liquid; iodine in crystals is 
of a dark color. Its vapor is violet. 



SOME FAMILIAR COMPOUNDS USED IN MEDICINE 189 

Emetics. — Emetics are substances which produce vomiting. 
Among the cathartics are mentioned calomel, or mercurous 
chloride, magnesium sulphate, magnesium citrate, and milk 
of magnesia. 
The caustics include sodium and potassium hydroxides, silver 
nitrate, and zinc chloride. 

Antidotes. — Antidotes are substances used to counteract poisons. 
They may do this by destroying the nature of the poison, 
as acids neutralize bases, or as magnesium sulphate renders 
carbolic acid harmless; or they may form insoluble com- 
pounds with the poison, as white of egg coagulates with 
mercuric chloride; or they may counteract the poison by 
their physiological action, as atropine counteracts morphine. 



REVIEW QUESTIONS 

1. Why has the discovery of bacteria changed medical 
science ? 

2. How does the use of hydrogen peroxide prevent blood 
poisoning 1 ? 

3. Why could not sodium peroxide be used in place of hy- 
drogen peroxide? 

4. How do you explain the frothing when dioxogen is poured 
upon a sore? 

5. How would you make formic acid from wood alcohol? 

6. How does carbolic acid differ from phenol? 

7. What are the various names given to the compound hav- 
ing the formula HgCl 9 ? 

8. How could sodium sulphite be made as a by-product in 
the reduction of mercury from cinnabar? 

9. What is the valence of mercury in the compound 
Hg(N0 3 ) 2 ? 

10. How can you use the formation of iodoform as a test for 
ethyl alcohol? 

11. What objections can you mention to the use of sulphur 
dioxide to disinfect rooms? 

12. How is boric acid made from colemanite? 



190 A PRACTICAL CHEMISTRY 

13. What is the chemical name of chloroform? 

14. What is bleaching powder? 

15. How does the manufacture of chloroform differ from the 
manufacture of iodoform? 

16. Ethyl ether used to be called si^lphurie ether; how did it 
get this name? 

17. What compounds are formed during the manufacture of 
ether? 

18. Why did Davy call nitrous oxide laughing gas? 

19. In what sort of operations can cocaine be used? 

20. State the injurious effects of nicotine.- 

21. Compare the properties of iodine with those of chlorine. 

22. What compound resembles bromine vapor in color? 

23. Under what circumstances are emetics necessary? 

24. What are some of the best emetics? 

25. How could you tell calomel from corrosive sublimate? 

26. What antidote could you give for mercurous nitrate? 

27. What chemical action might take place in the stomach if 
one took milk of magnesia and drank lemonade? 

28. How can you make silver nitrate? 

29. Discuss the different ways in which antidotes act. 

30. Why not give oil to a child that has eaten matches? 

31. What treatment would you employ in a case of poisoning 
with Paris green? 



CHAPTER XVII 
FIBERS 

In our study of things pertaining to our bodily ne- 
cessities and comforts we want to give some thought to 
the foundation materials of which clothing, carpets, tapes- 
tries, rope, and the like are made. We must learn some- 
thing of the processes by which they are obtained and the 
sources from which they come, the methods of removing 
undesirable colors from them and dyeing them to the 
desired hues. The chemistry of these processes, although 
in many cases quite complicated, is often not so extensive 
as the mechanical operations involved. These foundation 
materials may be classed under the general term " fibers/ ' 
These, in turn, are subdivided into those derived from the 
vegetable kingdom and those from the animal. In the 
former class we have cotton, flax, hemp, jute, china grass, 
manila and sisal, and artificial silk. Among animal fibers 
are wool and silk. 

In all these vegetable fibers there is but one important 
compound, cellulose, to which the formula C 6 II 10 O 5 is as- 
cribed, but its real structure is no doubt very complicated. 
While the composition of cellulose is apparently the same 
in all these various fibers and the percentage of other 
constituents is comparatively small, yet in appearance and 
properties they are all different. 

Surrounding the seed in the cotton bole is the soft 
downlike substance which we call cotton wool. This is 
picked by hand, and the seeds are removed by passing it 

191 



FIBERS 



193 



through a cotton gin. Referring to your history, you will 
recall the invention of the cotton gin by Eli Whitney 
and the great saving of labor brought about by its use. 
"Whitney's gin consisted of revolving circular saws, 
through which the cotton passed. Modifications of his 
machine are still used, but roller gins are also much used. 
After this it is spun into thread and woven into cloth 




Fig. 72. — A Cotton Gin. 



of many different kinds. The goods (or the thread before 
weaving) must be bleached. This is done by means of 
chlorine, as described in a previous chapter. 

Linen. — From the under bark of flax a fiber is obtained 
which is woven into cloth and called linen. Linen thread 
and cord also find a great variety of uses. 

The growing of flax and the preparation of the fiber 
for weaving require much labor and involve many opera- 



194 A PRACTICAL CHEMISTRY 

tions. After the plants are grown they are pulled up by 
the roots and the seed removed by means of large iron 
combs. The stalks are next subjected to the process of 
retting, by which the woody inner portions are partially 
rotted. This is accomplished either by soaking the flax 
in water or spreading it on the grass in the dew. No 
doubt from the chemical side the process of retting is 
quite complicated, as it is a form of fermentation brought 
about by a particular sort of bacteria. A method of 
chemical retting by the use of sulphuric and hydrochloric 
acids is being used now to some extent. 

After retting the flax is next broken by passing it 
through corrugated rollers, which break to pieces the 
woody stalks. The fibers must next be scutched, or cut 
lengthwise by numerous knives on a revolving wheel, 
after which it is hackled or combed through numerous 
iron combs, which separate the long threads from the 
short tow. The flax is now ready for spinning into thread 
and weaving. The bleaching of linen is accomplished in 
much the same as for cotton, i. e., by means of calcium 
hypochlorite or bleaching powder. Much grass bleaching 
is also done by spreading the linen on the grass in the 
sunshine and wetting it from time to time. 

Hemp. — Hemp is another plant somewhat similar to 
flax in that its bark by similar processes may be used as 
fiber, though usually for coarser textiles. 

Manila Hemp. — Manila hemp, or, as it is usually called, 
manila, is a sort of banana found in the Philippines. The 
outer portion of the fiber is our most important cordage 
material, while the inner part is made into many different 
kinds of clothing. 

Another entirely different sort of hemp is sisal. It is 
grown in Mexico, Central America, and Florida. It is 
also an important rope fiber, though not so valuable as 
manila. 




PULLING FLAX. 




harvesting hemp. 
Fig. 73. — The Fiber Industry. 



196 



A PRACTICAL CHEMISTRY 



A number of other plants used for fiber pass under 
the general name of hemp, but it will be noticed from 
what has been said that this classification does not depend 
upon the family to which they belong, but rather upon 
the character of the fiber. Another plant used for coarse 
textiles is called jute. Its fiber is obtained in a manner 
similar to that of flax. It is used for coat linings, cordage, 
and bagging. 

China-Grass or Ramie. — China-grass is a plant some- 
thing like a nettle which furnishes a beautiful silklike 
fiber which can be used for a great variety of fabrics in 
place of cotton. It is also used in making mantles for 
gas lights (see p. 89). Its use is quite modern. 

Artificial Silk. — One of our modern industries is the 
manufacture from cellulose of a silk which rivals the 
natural article. This is accomplished by chemical reac- 
tions which are too 
complicated to be 
discussed here, but 
are fast revolution- 
izing the silk busi- 
ness. The fiber is 
as beautiful as real 
silk, and many can- 
not tell the differ- 
ence between them. 
Natural silk is ob- 
tained from the silkworm. After the worm has grown 
to its full size it begins to spin a cocoon of one continuous 
double thread of silk from 800 to 1,000 yards in length. 
The worm is about four weeks old when it begins to spin, 
and the spinning requires about three days. During 
the chrysalis period the cocoons are sorted, those in- 
tended for breeding are put aside, while the others are 
killed in some convenient way as by placing in a steam 




Fig. 74. — Silk Cocoons. About M natural size. 



FIBERS 



197 



bath or in a hot oven. The cocoons are soaked in warm 
water to remove the gummy substance and the threads 
from several put together as one thread and reeled into 
skeins. After this the silk is cleaned, doubled and 
twisted, scoured, glossed, sometimes bleached, dyed, and 
weighted. The bleaching is done mostly with sulphur 
dioxide made by burning sulphur. The silk is thoroughly 
washed with the aid of dilute hydrochloric acid, and then 
with water containing some sodium carbonate, and finally 




Fig. 75. — A Spinning Machine. 



with water. The wet silk is next hung up in closed cham- 
bers and subjected to the sulphur dioxide fumes. Silk is 
also bleached by steeping with hydrogen peroxide solution 
containing a little ammonia. 

The weighting of silk is intended to make it heavier 
by adding to it metallic salts. This process adds nothing 
to the intrinsic value of the silk and destroys its dura- 
bility. 

Wool. — The fleece as it comes from the sheep contains 
a considerable amount of fatty matter called suint or 
yolk, as well as burrs and sand. The suint is removed by 



198 A PRACTICAL CHEMISTRY 

washing in soft water with soap. As previously men- 
tioned, the suint is used as a source of potassium car- 
bonate. The burrs and sand are removed by mechanical 
pickers when possible, but very small burrs are subjected 
to carbonization. This is accomplished by first putting 
the wool into a salt solution, as aluminium or magnesium 
chlorides, or into acid vapor and then baking it. The 
vegetable matter is thus carbonized and, after the wool is 
thoroughly washed and dried, falls out as a black powder. 

After properly mixing the wool it is slightly oiled and 
carded, spun, and woven. 

The bleaching of wool is accomplished in much the 
same way as in the case of silk, by means of sulphur 
dioxide. 

Dyeing. — From almost the earliest historical times it 
has been customary to change the color of fabrics by 
means of dyes. Until comparatively recent years dyes 
were derived from animal, mineral, and vegetable sub- 
stances, but at present a very large number of our most 
valuable dyes are made synthetically, starting with sub- 
stances found in coal tar, and are on this account called 
coal tar colors. The substances obtained from the coal tar 
are mostly hydrocarbons, among which are benzene, 
toluene, xylene, naphthalene, and anthracene. The chemi- 
cal reactions by which these are converted into dyes are 
quite complicated and involve a considerable knowledge 
of the compounds of carbon. Other dyes which were for 
many years obtained from natural sources are now being 
made synthetically in large quantities. 

Dyes are of two sorts : those which are permanent when 
used alone without a mordant to cause them to stick to 
the yarn, and those which require a mordant. A number 
of dyes which require a mordant for vegetable fabrics do 
not require one for wool and silk. 

In most cases the yarn is dyed before weaving, but some- 



FIBERS 199 

times cloth is dyed. It is always necessary that the ma- 
terial be free from grease and dirt and in many cases it 
must be bleached before dyeing. When the dyes do not 
need a mordant the goods are simply boiled with a solu- 
tion of the dye, washed, and dried. 

The object of a mordant is to cause the dye to stick to 
the yarn, and for this purpose compounds which form in- 
soluble substances with the dye are used. Very often 
the dye when combined with the mordant is of an entirely 
different color from what it is when alone. Thus logwood 
is brown, but mordanted with ferric chloride it is black. 

Among the most common mordants are a number of 
metallic salts, such as alum, ferric salts, and salts of tin. 
The goods or yarn after being saturated with a solution of 
a mordant of this kind is treated next with some chemical, 
such as ammonia or a phosphate, which will render the 
mordant insoluble. The yarn is after this boiled with the 
dye, washed, and dried. The combination of mordant and 
dye is spoken of as a lake. The student may easily illus- 
trate the formation of a lake by mixing some alum and 
litmus solutions and then adding ammonium hydroxide. 
He will observe that the precipitate of aluminium hydrox- 
ide thus formed is colored blue by the litmus and the dye 
is no longer in solution, but has formed an insoluble 
compound with aluminium hydroxide. 

Mercerizing. — We often hear of mercerized cotton goods 
and we have observed that such goods have a silky finish. 
This is accomplished by stretching the cotton yarn or 
fabric tightly over a frame and then treating it with a 
solution of sodium or potassium hydroxide. The alkali 
renders the cotton soft and causes it to shrink, but since 
it is held tightly by the frame the result is a stretching of 
the fiber which gives it a glossy appearance. 

The process is named from John Mercer, who discovered 
the action of alkalies on cotton. The process was not 



200 A PRACTICAL CHEMISTRY 

made practical, however, till about forty years after his 
discovery. 

SUMMARY 

Vegetable fibers include cotton, flax, hemp, jute, china grass, 

manila, sisal, and artificial silk. In all of these cellulose is 

the most important compound. 
Animal fibers are wool and silk. 
Cotton is a wool-like substance which surrounds the seeds of 

the cotton plant. The seeds are removed by ginning. 
Flax, hemp, sisal, manila and jute are plants which furnish 

fibrous bark which may be separated from the woody stems 

by retting, breaking, and scutching. 
Sisal and manila are used for cordage. Jute and hemp are 

used in making coarse textiles. The fabric made from flax 

is called linen. 
China grass is a silklike fiber from a plant which resembles a 

nettle. 
Artificial silk is made from wood by complicated chemical 

processes. 
Silk is made by silkworms, each worm furnishing from 800 

to 1,000 yards of thread. The silk thread is cleaned, twisted, 

doubled, scoured, glossed, bleached, dyed. 
Cotton is bleached with chlorine, wool and silk with sulphur 

dioxide, while flax is often bleached on the grass. 
Dyeing. — Dyeing is done both with and without a mordant. A 

mordant is a substance used to cause the dye to stick to the 

cloth or yarn. Many dyes are made from substances found 

in coal tar. 
A lake is a combination of a mordant and a dye. 
Mercerizing. — Mercerizing is accomplished by treating the cotton 

goods with a solution of sodium or potassium hydroxide while 

it is tightly stretched. This gives the goods a silky finish. 



FIBERS 201 

REVIEW QUESTIONS 

1. To which of the sources of fibers are we indebted for the 
greater part of our clothing? 

2. If the cellulose in the various sorts of vegetable fibers is 
the same compound, how do you account for the great differences 
between them? 

3. How does cotton differ from wool? 

4. What is meant by grass bleaching? 

5. What part do bacteria take in the preparation of linen? 

6. What is included under the name hemp? 

7. What fiber may sometime supplant cotton? 

8. How is the silk business being changed? 

9. Why is it necessary to kill the silkworms in the cocoons? 

10. Why would silk be better without weighting? 

11. How is it possible to make potassium carbonate from 
suint? 

12. Why are the coal-tar colors so named? 

13. How is the dye industry being changed? 

14. How can a dye be removed from solution? 

15. How do the methods of dyeing animal fibers often differ 
from those used in dyeing cotton? 

16. What is gained by mercerizing cotton? 



CHAPTER XVIII 
CLEANING MATERIALS 

Some one has defined dirt as " matter but of place." 
Whether this definition be accepted or not it is certain 
that in all our cleaning processes we attempt to remove 
matter objectionable because of its presence. 

In all cleaning a few simple principles are involved. 
No one would think of removing paint from clothing with 
cold water, or of using gasoline to wash candy from the 
face of a child. In other words, a suitable solvent is 
sought for each particular form of matter — gasoline for 
paint and water for candy. Again, everyone appre- 
ciates the importance of mechanical assistance in the re- 
moval of dirt, hence the addition of sand to soap. A third 
principle less often considered is the softening of the sub- 
stance to be cleaned. Clothing is soaked in water for 
some hours previous to washing for this purpose. At the 
same time we attempt to soften the dirt. This brings us 
to the consideration of substances used as solvents, and 
of those used to loosen the dirt by softening it and by 
softening the material from which it is to be removed. 
The second of these includes all sorts of soaps, soap 
powders, and washing compounds. 

Soap. — In the chapter on foods we sought to explain 
the composition of the natural fats ; it will be recalled 
that these compounds were defined as the glyceryl salts 
of complicated organic acids, and, moreover, that the 
glyceryl radical is derived from the familiar substance, 

202 



CLEANING MATERIALS 



203 



glycerine. We must now learn that fats, when boiled 
with sodium hydroxide, undergo decomposition, the so- 
dium taking the place of the glyceryl radical and forming 
sodium salts of the acids present. These sodium salts of 
fatty acids are called soaps. Potassium hydroxide pro- 
duces the same reaction except that the soaps formed are 




Fig. 76. — Cutting Large Blocks of Soap. The soap when taken from the frames 
is cut into slabs by pushing it against tightly stretched wires. In the picture 
these slabs are being cut into small blocks. Reproduced by the courtesy of Colgate 
and Co. 



soft and do not solidify, while the sodium soaps are hard. 
All of our ordinary soaps are sodium soaps. 

The process of decomposing a fat by a hydroxide is 
called saponification. During saponification the glyceryl 
radical unites with hydrogen and oxygen from the hy- 
droxide and forms glycerine. 

All are familiar with the great number of soaps on the 



204 



A PRACTICAL CHEMISTRY 



market. The number is almost limitless because not only 
can we have a soap to correspond to nearly every natural 
fat, but we may use mixtures of fats. Then, to compli- 
cate matters more, various foreign substances are added 
to the soap, such as sand, borax, rosin, sulphur, tar, and 
numerous other things, to say nothing of perfumes and 
colors. Some of these things are desirable in soaps for 
special purposes, but for bathing the purest soap is the 
best. Castile soap, which has long been regarded as a 




Fig. 77. — Various Processes in the Manufacture of Shaving Soap. Reproduced 
by the courtesy of Colgate and Co. 

standard, is made from olive oil. Palm oil gives a fine 
white soap. Much of our common soap is made from tal- 
low. 

The manufacture of soap may be described somewhat 
as follows: The fat is melted by means of steam and 
run into large kettles of ''lye" (sodium hydroxide solu- 
tion). The mixture is next boiled until saponification is 
complete. At this stage a quantity of common salt is 
added, which causes the soap to come to the top. Be- 
neath is a mixture of brine, alkali, and glycerine. This 



CLEANING MATERIALS 205 

mixture is drawn off, neutralized, filtered, and evapor- 
ated in vacuum kettles, the salt crystallizes out, and 
glycerine is obtained. The soap is washed and mixed 
with any of the numerous substances put into soaps and 
cooled in large "frames," cut into blocks, stamped and 
dried. Soap is made to float by blowing air into it while 
it is being melted. 

Washing compounds are usually mixtures of soap with 
such substances as borax, sodium carbonate, sodium hy- 
droxide, et cetera. They are, no doubt, destructive to 
clothing and injurious to the hands of the laundress. 

Action of Soap. — It has already been stated that soaps 
are intended to soften the dirt and also the object to be 
cleansed. In addition to this, soap emulsifies the fats and 
other dirt so that they float away mechanically in the 
water. When the soap contains sand, it serves mechan- 
ically to scratch away the dirt. 

Hard and Soft Water. — While many solvents find place 
in cleaning, yet the common one is water. On account of 
its great solvent power water is seldom pure. For wash- 
ing, however, the question is not so much whether the 
water is pure as whether it contains anything which will 
render soap useless by changing it into an insoluble com- 
pound. A number of different salts — such as acid calcium 
carbonate, Ca(HC0 3 ) 2 ; calcium sulphate, CaS0 4 ; mag- 
nesium sulphate, MgS0 4 — and others often found in water 
will do this. Water, then, which contains anything that 
will render soap insoluble and thus interfere with its 
forming a lather is said to be hard water. On the other 
hand, water that lathers well is soft. 

Since the action of hard water on soap is explained by 
saying that the calcium or other metal of the dissolved 
salt in the hard water takes the place of the sodium in 
the soap and forms an insoluble soap, it follows that if an 
excess of soap is added to hard water it will remove all 



206 A PRACTICAL CHEMISTRY 

of the calcium or other metal and render the water soft. 
The excess of soap may then form a lather. Other meth- 
ods of softening water exist. If the water contains only 
acid calcium carbonate, it may be softened simply by 
boiling, since this compound decomposes at the boiling 
temperature into carbon dioxide and water and calcium 
carbonate. The calcium carbonate being insoluble can- 
not react with the soap. Water that can be softened by 
boiling is called temporarily hard. Any water may be 
rendered soft by removing the mineral matter. This may 
be accomplished by the addition of sodium carbonate or 
of borax after the temporary hardness has been removed 
by boiling. These compounds, sodium carbonate and 
borax, form with the calcium, magnesium, and other salts 
in the water insoluble carbonates and borates and leave 
sodium salts (which cannot affect the soap) in solution 
in the water. 

When studying baking powders we learned about 
sodium bicarbonate and how it is made ; now sodium car- 
bonate, Na 2 C0 3 , is made from the bicarbonate by heating 
it, thus 2NaHC0 8 = Na 2 C0 3 + H,0 + C0 2 . Hence in 
the Solvay process the product is heated if sodium car- 
bonate is desired, and the CO, obtained is used over again. 
Sodium carbonate is also made by the Leblanc process. 
The first step in this process is the same as the manu- 
facture of hydrochloric acid. Common salt and sulphuric 
acid are heated together until the reaction 

2NaCl + H 2 S0 4 = 2HC1 + Na 2 S0 4 
(Sodium (Hydrochloric (Sodium 

chloride) acid) sulphate) 

is complete. The hydrochloric acid passes off as a gas 
and goes into a large condensing tower filled with coke 
over which water is flowing. The water dissolves the 
acid, forming a strong solution of hydrochloric acid, which 



CLEANING MATERIALS 207 

collects in tanks at the bottom. This solution is further 
concentrated by letting more of the gas pass over it on 
its way to the condensing tower. The solution of hydro- 
chloric acid thus formed is the ordinary hydrochloric acid 
of commerce. 

The sodium sulphate made in this first step is brought 
to a high temperature to remove any sulphuric acid that 
may be clinging to it, and is then mixed with coke and 
limestone. The mixture is thrown into a furnace, where 
two reactions take place. In the first of these the sul- 
phate is reduced by the hot carbon to the sulphide : 

Na 2 S0 4 + 2C = Na 2 S + 2C0 2 

(Sodium 
sulphide) 

The sodium sulphide next reacts with the limestone or 
calcium carbonate thus: 



3 



Na 2 S + CaC0 3 = CaS + Na 2 C0 2 
(Sodium (Calcium (Calcium (Sodium 
sulphide) carbonate) sulphide) carbonate) 

The sodium carbonate is dissolved out with water, the 
solution filtered and evaporated. The carbonate crystal- 
lizes out with ten molecules of water of crystallization, 
Na 2 C0 3 + 10H 2 O. 

Just as acid sodium carbonate is called baking soda, so 
sodium carbonate is called washing soda. 

"When washing soda is added to hard water contain- 
ing mineral salts, as, for example, calcium sulphate, in- 
soluble carbonates are formed and thus the mineral is 
removed and the water made soft. 

Among other solvents used in what is known as dry 
cleaning, gasoline plays an important part, particularly 
where woolen goods are concerned. All grease and fatty 
substances dissolve readily in gasoline, and most of the 



208 A PRACTICAL CHEMISTRY 

other dirt either dissolves or floats out mechanically. 
Benzine is of much the same character and is used in the 
same way. Both are derived from petroleum by frac- 
tional distillation, and both are. highly inflammable sub- 
stances which must not be used in the same room with 
fire. 

Carbon Tetrachloride. — Carbon tetrachloride, CC1 4 , a 
compound closely related to chloroform but having no 
hydrogen and one more chlorine atom, is taking the place 
of gasolene and benzine for this work", since it is not 
combustible. Carbon tetrachloride does not have the 
anaesthetic properties of chloroform ; its odor also, though 
not unpleasant, is quite different. Carbon tetrachloride 
may be made by the action of chlorine on carbon bisul- 
phide. 

Most of the other cleaning substances depend largely 
upon chemical action for their efficiency, although in 
many cases the chemical changes are quite complicated. 
Thus we have the use of ammonia to remove grease and 
other stains. Oxalic acid, C 2 H 2 4 , a white crystalline acid, 
salts of which occur in a number of plants, is also used 
to remove many stains. It is particularly good for the 
removal of ink. Oxalic acid is a deadly poison and care 
should be employed in handling it. Such stains as iron 
rust may also be removed with oxalic or with hydrochloric 
acid. In some laundries bleaching powder finds a place 
in the washing processes. 

It is needless to say that all of these and other strong 
chemical substances, which are being used so freely in 
modern laundries, are destructive to clothing and should 
be condemned. Pure soaps, steam, and hot water should 
be restored to the larger place which they deserve in 
laundry work. 

Rust and corrosion are removed from metals by means 
of acids. Hydrochloric and sulphuric acids are both used 



h CLEANING MATERIALS 209 

for steel, while oxalic acid is used for cleaning brass. 
When steel is greasy it must first be burned off, it is then 
soaked in dilute acid, scoured, and then washed with 
water containing washing soda. 

The sand is removed from castings sometimes with 
hydrofluoric acid, a substance having the formula HF. 
It is a very volatile liquid and is used in water solution. 
Hydrofluoric acid is remarkable for its activity, it attacks 
the sand (silicon dioxide), converting the silicon into a 
fluoride and the oxygen into water. Thus, 

Si0 2 + 4HF = SiF 4 + 2H 2 
(Silicon (Silicon 

dioxide) tetrafluoride) 

It also will attack glass and almost everything else. 

Hydrofluoric acid is made by the action of sulphuric 
acid on fluorspar, a mineral composed of calcium fluoride, 
CaF 2 . The acid is volatilized and collected in water. The 
acid and all of its soluble salts (the fluorides) are violent 
poisons. In contact with the skin or any of the mem- 
branes of the body the free acid produces ugly ulcers. 
The equation representing the manufacture of the acid is : 

CaF 2 + H2SO4 = CaS0 4 + 2HF 
(Calcium , (Calcium (Hydroflu- 

fluoride) sulphate) oric acid) 



SUMMARY 

Cleaning. — Cleaning, which consists in removing undesirable mat- 
ter, involves three things — a suitable solvent, mechanical agi- 
tation, and substances to soften both the dirt and the ma- 
terial from which it is to be removed. 

Soaps. — Soaps are usually sodium salts of fatty acids, and are 
made by boiling fats with sodium hydroxide. If potassium 



210 A PRACTICAL CHEMISTRY 

hydroxide is used, a soft soap is obtained.- A soap may con- 
tain the sodium salt of only one fatty acid, or it may con- 
tain many such salts. It may also contain many different 
kinds of foreign matter, such as sand, sulphur or borax. 

Saponification is the decomposition of a fat by means of a 
base. Soap softens the dirt and also the material to be 
cleansed. It also emulsifies fats and other dirt so that they 
may float away in the water. 

Water is hard when it contains mineral salts which con- 
vert soap into an insoluble compound incapable of producing 
a lather. Among the most common of these mineral salts 
are acid calcium carbonate, Ca(HC0 3 ) 2 , calcium sulphate, 
CaS0 4 , and magnesium sulphate, MgS0 4 . Acid carbonates 
may be decomposed and removed by boiling, hence waters 
containing them are said to be temporarily hard. Other 
salts may be removed by converting them into insoluble 
compounds, as by the addition of sodium carbonate, called 
washing soda. Washing soda is made by heating sodium 
bicarbonate from the Solvay process or by the Leblane 
process. 
Dry Cleaning. — Dry cleaning is done with benzine, gasoline, and 
carbon tetrachloride. 

Oxalic acid is used to remove stains and to clean brass; sul- 
phuric and hydrochloric acids to clean steel, and hydrofluoric 
acid to remove sand from casting. Hydrofluoric acid and its 
salts are dreadful poisons. The acid produces ugly ulcers. 
It is very active and will etch glass. It is made from cal- 
cium fluoride and sulphuric acid. 

REVIEW QUESTIONS 

1. Why is hot water better than cold for washing clothing 1 ? 

2. What is the composition of a soap made from stearin and 
potassium hydroxide? 

3. What by-product would be formed in the manufacture of 
this soap 1 ? 

4. How would this soap differ from one made from stearin 
and sodium hydroxide? 



_L 



CLEANING MATERIALS 211 

5. Why is it possible to have so many kinds of soaps'? 

6. Explain the formation of calcium carbonate in kettles in 
which water is boiled. 

7. If potassium carbonate were mixed with soap, what effect 
could it have upon temporarily hard water? 

8. What upon other kinds of hard water? 

9. Why do bubbles of carbon dioxide escape from tem- 
porarily hard water while it is being heated? 

10. What by-products are formed in the Leblanc process? 

11. Which is worth more, a pound of freshly made crystals 
of washing soda or an equal weight of washing soda that has 
been exposed to the air for some time? 

12. Write the chemical equations representing the various 
steps of the Leblanc process. 

13. Why is carbon tetrachloride less dangerous than gasoline 
for dry cleaning? 

14. What are the objections to modern laundry methods? 

15. How would you clean corroded brass? 

16. How could you etch your name on glass? 

17. What danger is involved in this process? 



CHAPTER XIX 
BUILDING MATERIALS 

As man has slowly progressed in civilization the ma- 
terials with which he has done his building have under- 
gone many changes, but notwithstanding all these changes 
the earth is, and ever has been, the source of the most 
substantial products for this purpose. 

Silicon and Silicates. — The student has perhaps ere this 
grasped the idea that in organized nature — in plants and 
animals — carbon is the central element; likewise in the 
earth another element, not very unlike carbon in its 
chemical behavior, prevails. This element is silicon. 

Simplest of its compounds is the white sand of the 
seashore, chemically known as silicon dioxide, Si0 2 , and 
familiar to all. The same compound occurs also as quartz, 
flint, agate, rock crystal, and many other forms of less 
importance. While these are all silicon dioxide, they 
receive their various names from differences of color due 
to small traces of other substances, and from differences 
in crystal structure. In the amethyst we have an ex- 
ample of crystallized quartz colored with a little man- 
ganese, while in quartz rock we often find quartz massive 
and white. Nearly all forms of silicon dioxide find uses 
in building and manufacture. 

Chemically speaking, silicon dioxide is rather inactive ; 
it may, however, be brought into combination by heating 
with an alkaline hydroxide, as sodium hydroxide, or by 
fusion with sodium or potassium carbonate. 

212 



BUILDING MATERIALS 



213 



We will carefully examine the reaction with sodium 
carbonate, since it is of interest to us, both on account of 
the "product formed and also because it introduces us to 




Fig. 78. — Building a High School, 



a new class of compounds. The reaction with sodium 
carbonate is represented thus: 

Na 2 C0 3 + Si0 2 = Na 2 Si0 3 + C0 2 
(Sodium 
silicate) 

A mixture of sodium carbonate and potassium carbon- 
ate is often heated with silicon dioxide until a reaction 
similar to the one just given has been completed. The 
product is called water glass and is a mixture of sodium 



214 



A PRACTICAL CHEMISTRY 



and potassium silicates. Its name is derived from the 
fact that the substance is a sort of glass, but, unlike glass, 
is soluble in water. 

Looking again at this equation it will be seen that car- 
bon dioxide of the carbonate is replaced by silicon dioxide 
with the formation of a silicate. These silicates are salts 
of a silicic acid of the formula, H 2 Si0 3 . This is but one of 




Fig. 79. — A Marble Quarry in Vermont. The marble is removed in large blocks 
which are afterwards cut and polished. 



a number of acids of silicon very complicated in structure 
from which almost numberless silicates are derived. 
Many of these silicates have definite mineral names. The 
rocks are made up largely of mixtures of these various 
minerals and quartz cemented together by silicates. Gran- 
ite, which contains quartz, feldspar, and mica, is an illus- 
tration of this. 

Gf silicon as a free element we need to say little. It 
does not thus occur in nature, and science thus far has 
found but few uses for it. It should be mentioned, how- 



BUILDING MATERIALS 215 

ever, that, like carbon, silicon occurs in three modifica- 
tions: amorphous, crystalline, and graphitoidal. 

Carbonate Rocks. — In some sections, however, we find 
another class of rocks well represented by limestone or 
calcium carbonate. In fact, so far as building purposes 
are concerned, calcium carbonate is the only one of these 
compounds to be considered. But calcium carbonate oc- 
curs in as many forms as silicon dioxide. Among these 
may be mentioned, besides limestone, which is impure, cal- 
cium carbonate, marble, which is a much purer and more 
or less crystalline form of limestone ; calcite, a crystalline 
variety of considerable purity; and Iceland spar. The 
last mentioned is pure calcium carbonate in transparent 
crystals which have the power of double refraction, i. e., 
they bend the light in two directions at once, so that ob- 
jects seen through them appear double. 

Calcium carbonate also occurs in the shells and the 
bony structure of many animals. Coral is an illustration 
of rock made up of such structures which were once the 
skeletons of coral polyps. Chalk, likewise, is a mass of 
shells of microscopic animals which have slowly accumu- 
lated at the bottom of the sea. These little animals live 
and die in the water, and their shells form the ooze on 
the ocean's bottom. The chalk deposits are consolidated 
ooze. 

Under hard water we learned that acid calcium car- 
bonate is soluble in water, while calcium carbonate is in- 
soluble. In other words, calcium carbonate will dissolve 
in water containing carbonic acid or carbon dioxide. The 
caves which are found in limestone countries are formed 
by the dissolving out of the limestone by water containing 
carbonic acid. Carbon dioxide taken up by the water 
from the air forms carbonic acid, which in turn unites 
with calcium carbonate, forming acid calcium carbonate. 
These reactions may be expressed thus : 



216 A PRACTICAL CHEMISTRY 

C0 2 + H 2 = H 2 C0 3 
and 

H 2 C0 3 + CaC0 3 = H 2 Ca(C0 3 ) 2 

Lime. — It has been mentioned that acid calcium car- 
bonate decomposes at the temperature of boiling water. 
Limestone requires a much higher temperature. The 
stone is broken up, put into a kiln and heated for some 
hours with a coal or wood fire. Quicklime, or calcium 
oxide, CaO, is the product. 

CaC0 3 = CaO -f- C0 2 ; the carbon dioxide escapes. Cal- 
cium oxide, like many other oxides, reacts with water, 
forming a basic compound, calcium hydroxide, or slaked 
lime. 

CaO + H 2 = Ca(OH) 2 
(Calcium 
hydroxide) 

This process is called slaking. In slaking lime much 
heat is produced. The reaction proceeds slowly at first 
but soon becomes quite violent. The water should be 
added in small quantities at first, and, after the reaction 
has advanced, in much larger quantities. 

Calcium hydroxide in solution forms lime water, much 
used as a reagent in testing for carbon dioxide. It is also 
used in medicine. When mixed with water to about the 
consistency of milk calcium hydroxide is called milk of 
lime or whitewash, and is used in place of paint. 

Clays. — Rocks gradually disintegrate in the weather 
and crumble into dust. Different kinds of rocks produce 
different kinds of soils. Thus the weathering of sand- 
stone gives us sandy soils, and the weathering of feld- 
spar produces clay. Feldspar contains silicates of the 
alkalies, sodium, and potassium, and of aluminium. Dur- 
ing the weathering, as the rock disintegrates, the alkali 
silicates wash away, being soluble, but the aluminium sili- 



BUILDING MATERIALS 



217 




cate remains. This impure aluminium silicate is called 
clay. Admixtures of other things produce many different 
clays, most of which are useful for some particular pur- 
pose. 

Masonry. — We have thus enumerated the materials re- 
quired by the mason in building walls. The granite, mar- 
ble, and sandstones are quarried from 
the rocks of the hills ; bricks are made 
from the clay by pressing it into shape, 
drying the bricks thus formed, and 
baking them in large kilns. Then be- 
tween the stone or the bricks must be 
put mortar or cement, and the walls 
must be covered with plaster. 

Mortar. — Mortar is a mixture of lime 
and sand. The lime is slaked as de- 
scribed above and mixed with water 
till a mixture about like cream is 
formed. Sand is stirred into this until 
the mortar is thick enough to spread 
well. The hardening of the mortar 
depends, first, upon the drying out of 
the water, and, second (and this is of 
greater importance), upon the absorp- 
tion of carbon dioxide from the air changing the lime back 
into limestone. Thus 

Ca(0H) 2 + C0 2 = CaC0 3 + H 2 

The latter process goes on slowly for a long while, the 
mortar gradually becoming harder. The sand serves to 
make the mortar porous and also adds to its bulk. 

Plaster. — Plaster is much the same in composition as 
mortar, though richer in lime. Hair also is added to make 
it hold together more firmly. Frequently, in addition to 
these, plaster of Paris is mixed with plaster. Plaster of 



Fig. 80. — D iagram 
Showing the Shape 
of Perfect Gypsum 
Crystals. They are 
nearly transparent 
and are easily 
broken. When their 
water of crystalliza- 
tion is driven out 
they fall to a white 
powder. 



218 A PRACTICAL CHEMISTRY 

Paris is calcium sulphate. The mineral from which it is 
made is gypsum. Plaster of Paris does not harden by ab- 
sorbing carbon dioxide as mortar does; the process may 
be explained as follows: Gypsum as it comes from the 
ground is crystallized calcium sulphate of the formula 
CaS0 4 + 2H 2 0. When this is heated to about 115° C. 
part of the water passes off and the formula becomes 
(CaSOJ 2 .H 2 0, and the crystalline structure is destroyed. 
On being mixed with the right amount of water this 
powdered plaster of Paris again takes up water and forms 
a mass of microscopic crystals. Thus, 

(CaSOJ 2 .H 2 + 3H 2 = 2(CaS0 4 -.2H 2 0) 

Here we have a little hint of the importance of water 
of crystallization. Another illustration of a similar reac- 
tion of water is probably found in the hardening of ce- 
ment. 

Cements. — Cements are now much used in place of mor- 
tar. In many parts of the world large deposits of ce- 
ment rock are found. These vary in composition, but 
their composition may be illustrated by the following an- 
alysis of the cement rock of the Lehigh District, as given 
by the U. S. Geological Survey: 

Per cent. 

Silicon Dioxide, Si0 2 15.05 

Aluminium oride, A1 2 3 9.02 

Iron oxide, Fe 2 3 1.27 

Calcium carbonate, CaC0 3 . . 70.10 

Magnesium carbonate, MgC0 3 3.96 

The aluminium and iron are therefore present as silicates 
and the rock may be described as argillaceous limestone. 

Rock of this character is broken into small pieces, 
mixed with limestone, and then crushed. , This mixture 
is next dried and ground to a fine powder. It is then 






BUILDING MATERIALS 



219 



conveyed to the upper end of a long inclined rotary kiln ; 
as it slowly passes through this kiln it is gradually heated 




C£M£VT 




sS/f/V0 



5TO/V£ 



COA/C/?£T£ 



Fig. 81. — The Proportion in Which the Cement, Sand and Stone Should be 
Mixed in Making Concrete. 

to a temperature of about 1,300° C. and is converted 
into a "clinker." The clinker thus formed is mixed with 




Fig. 82. — Mixing Concrete by Hand. Although concrete when wanted in large 
quantities is usually mixed by machines, for small structures the work may be 
done by hand as shown in this picture. 



220 



A PRACTICAL CHEMISTRY 



a small amount of gypsum and again is ground to a fine 
powder. The cement is now ready for use in place of 
mortar or for making concrete. When used in place of 
mortar it is mixed with sand and water to the consist- 
ency of soft mud. The gypsum in the cement prevents its 

immediate hardening, 
but this process soon 
begins. The chemical 
changes produced by 
the water during the 
hardening of cement 
are not yet fully un- 
derstood. 

Concrete. — Masonry 
is fast giving place to 
concrete. This is due 
to the great stability 
of the latter and the 
ease with which it is 
made to conform to 
any desired shape. A 
wooden frame of the 
required form is first 
made {see Fig. 83) ; 
this is filled with a 
mixture of cement 
and crushed stone and 
sand {see Fig. 81). 
The cement soon hardens and the woodwork is removed. 
Concrete is often reinforced by steel rods and steel net- 
work imbedded in it. 

Metals Used in Building. — It is our purpose under this 
general heading to learn something of those metals which 
are more or less familiar to all. It would be next to im- 
possible to give an exact definition of the word " metal/ ' 




Fig. 83. — A Wood Frame in Which to Build a 
Concrete Wall. After the concrete has 
hardened the frame is removed. The cross 
wires keep the frame from spreading. 



BUILDING MATERIALS 221 

and yet it suggests a more or less heavy substance, having 
a peculiar luster, and capable of being polished. The 
metals usually are good conductors of heat and electricity. 
They may be hammered into sheets and drawn into wire, 
i. e., they are malleable and ductile. Again they are the 
elements which replace hydrogen in acids in the forma- 
tion of salts. The metals most used in building are iron — 
including its three forms — cast iron, steel, and wrought 
iron — lead, tin, zinc, copper, and aluminium. 

We shall hope to learn something of the properties of 
these, how they are extracted from their ores (i. e., the 
minerals which are worked to obtain them), and some- 
thing of their important compounds. Some of their com- 
pounds, however, will require further discussion under 
other headings. 

Iron. — The introduction of steel into building in place 
of wood has made possible the construction of the sky- 
scrapers and gigantic bridges of modern times. 

The story of steel is a long one and full of interest. 
It is but one of the many monuments which stand to the 
glory of chemical progress as a leading factor of our 
civilization. 

Iron, like other metals, is not taken from the earth in 
the free condition, but must be reduced from its ores at 
a very high temperature in blast furnaces. 

We may describe the blast furnace as a tower about 
90 feet high, widest at a point about 20 feet above the 
base, and tapering slowly toward the top and abruptly 
to the bottom. It is constructed of the most refractory 
firebrick on the inside and of iron plates without, riveted 
and banded together to withstand the heat and the great 
weight of the burning mass within. The bottom of the 
furnace is a fireclay hearth ; a few feet above this the 
furnace is penetrated by a number of blow-pipes called 
tuyeres, through which a hot blast is blown into the fur- 




Fig. 84. — The Woolworth Building. This highest office building in the world is 
built almost entirely of steel and concrete and is a monument to modern building 
materials. 



BUILDING MATERIALS 



223 



nace. The entrance to the furnace is at the top and is 
closed by a cone on to which are dumped the fuel and 
charge for the furnace. 
The cone lowers and per- 
mits the load to fall into 
the furnace and then closes 
automatically. From near 
the top of the furnace a 
huge pipe descends to the 
ground and conveys the 
gases from the furnace. 
These are used to heat 
great brickwork stoves, 
which in turn are employed 
to heat the air for the hot 
blast. Thus much heat 
which would otherwise be 
lost is returned to the fur- 
nace. 

The fuel of the blast 
furnace consists of coke, 
which is particularly 
adapted to this work on 
account of the intense heat 
which it produces and on 
account of its hardness, 
which enables it to with- 
stand the weight of the 
charge. The charge con- 
sists of the ore and flux. 
The ores of iron most used 
are hematite, Fe 2 3 ; limon- 
ite, 2Fe 2 3 .3H 2 ; magne- 
tite, Fe 3 4 ; and siderite, 
FeCOo. These exist in various modifications. 




Fig. 85. — Diagram of Cross-Section 
of a Blast Furnace. The mate- 
rials (fuel, ore and flux) are dropped 
into the furnace through the door, 
F, which is brought back to its place 
by the counterpoise, C. A is a large 
pipe which encircles the furnace and 
carries the blast which passes on 
through the tuyeres T, T (and oth- 
ers not shown) into the furnace. 
The course of the blast is indicated 
by the arrows. Most of the gases 
from the furnace escape through 
the "down comer," D, passing 
through a large dust box (not 
shown) and thence to the heating 
stoves. The slag runs out at E. 
The iron passes out through a hole 
just above the hearth, H. 



224 A PRACTICAL CHEMISTRY 

The flux must be of such a nature as to suit the ore; 
if the ore contains much silicon dioxide or other acid- 
forming substances the flux should contain a correspond- 
ing amount of base. In this case, limestone would be used, 
which would burn to lime in the furnace. On the other 
hand, should the ore contain basic substances, the flux 
must be correspondingly acid. The purpose of the flux 
is to unite with the impurities in the ore and form a 
glass-like slag which aids the movement of the material 
down through the furnace and at the same time protects 
the iron from oxidation. The furnace is fed by succes- 
sively dumping into it small car-loads of coke, flux, and 
ore through the opening at the top. The coke is burned to 
carbon monoxide, CO, and this, reacting with the oxygen 
of the ore, is converted into carbon dioxide, while the 
iron is reduced to the metallic state and is in a molten 
condition. 

The iron and slag fall to the bottom of the furnace and 
accumulate on the hearth, the slag floating on top of the 
iron and protecting it from oxidation. At regular inter- 
vals these are drawn out at openings which at other times 
are closed with fire clay. From the furnace the iron is 
conducted into numerous little channels in sand where 
it is permitted to cool and form little bars called "pigs." 
It is often, however, molded in a casting machine, or run 
into iron cars and poured while still nearly white hot into 
Bessemer converters, where it is made into steel. 

Pig Iron. — Pig iron, the product of the blast furnace, is 
cast iron. It contains, besides iron, silicon, carbon, sul- 
phur, phosphorus, and small quantities of numerous other 
substances. The silicon and carbon are the most impor- 
tant of these impurities. Several different kinds of cast 
iron are known. Two of these are white cast iron and 
gray cast iron. These colors depend upon how much of 
the carbon is combined with iron and how much is free 



BUILDING MATERIALS 225 

carbon or graphite — the more nearly all of the carbon is 
combined the lighter will be the color of the iron. 

Cast iron is the most impure of all forms of iron. It 
cannot be hammered into shape or welded. It is very 
brittle, but will sustain a great load vertically. Cast iron 
melts at a lower temperature than either steel or wrought 
iron (about 1,100° C). It is used to make castings by 
melting it and pouring it into molds in the sand. The 
student will notice numerous objects around him made of 
cast iron. 

Wrought Iron. — The purest iron used for commercial 
purposes is wrought iron, so named, probably, from the 
fact that it is worked into so many shapes during its 
manufacture. Not only is wrought iron unlike cast iron 
in composition but in nearly every other particular. It is 
tough and can, when hot, be bent into any shape or ham- 
mered out to any form. On this account it is the iron 
most used by the blacksmith. Its melting point runs as 
high as 2,000° C, and before melting it becomes pasty, so 
that two pieces can be welded or hammered together. It 
cannot be used for columns to support the weight of build- 
ings, but when made into wire or cables it is very strong. 

Wrought iron is made by burning out the impurities 
from cast iron. A mixture of cast iron and hematite is 
put into a form of reverberatory furnace known as a 
puddling furnace. The cast iron melts and its impurities 
unite with oxygen from the hematite and from the air. 
As the iron becomes more and more pure its melting point 
rises until it is no longer liquid. While in this pasty 
condition it is stirred or "puddled" and the more solid 
portions stick together and are removed in lumps. These 
lumps are squeezed, hammered, and rolled, until the slag 
is removed and the iron is in a tough fibrous condition. 

Steel. — It is next to impossible to define steel, as so 
many different alloys of iron made in so many different 



226 



A PRACTICAL CHEMISTRY 



ways bear this name. Formerly steel was made by heat- 
ing bars of wrought iron with charcoal in sealed iron 
boxes for days together. Some carbon thus enters the iron. 
Modern steel began to be made about 1856, when Besse- 
mer invented the converter. This is an oval crucible of 
several tons capacity, hung on trunnions, and so arranged 
that a blast of air may be forced in at the bottom. The 

molten pig iron from 
the blast furnace is 
poured into the con- 
verter and the blast 
forced through until 
all carbon and silicon 
are burned out. This 
point can be told by 
the color of the flame 
coming from the 
mouth of the con- 
verter. A quantity of 
a particular kind of 
cast iron known as 
spiegeleisen, sufficient 
to furnish the right 
amount of carbon to 
convert the whole 
charge into steel, is next added. Spiegeleisen contains 
carbon and manganese. It is previously carefully ana- 
lyzed so that the proper quantity may be added. The 
steel is then poured into large molds in which it hardens. 
The molds are removed while the ingots of steel are still 
quite hot. These are then dropped into a furnace where 
they can come to the same temperature throughout. The 
ingots are next passed through a succession of rolls which 
draw the steel out into long, tough fibrous bars and pre- 
vent its crystallizing while cooling. 




V»»»»W >»»»»»,S7777. 



Fig. 86. — The Bessemer Converter. The 
figure shows a vertical section of this appa- 
ratus. It consists of a large crucible of steel 
lined with fire-brick. The crucible is about 
15 ft. high by 8 ft. in diameter and is sup- 
ported by trunnions. An air-blast is 
forced through as shown by the arrows. 



.. 



BUILDING MATERIALS 227 



If the cast iron from which the steel is to be made 
contains sulphur and phosphorus, the Thomas-Gilchrist 
modification of the Bessemer process is employed. This 
consists in lining the converter with magnesium and cal- 
cium carbonates. Lime is also put into the converter 
with the iron. The sulphur and phosphorus become ox- 
ides which react with the lime, forming calcium sulphate 
and phosphate which float as a slag upon the top of the 
molten steel. This slag is used as fertilizer. 

A more expensive method of making steel is known as 
the open-hearth process. Cast iron and wrought iron and 
sometimes scrap iron are mixed and heated in a horizon- 
tal furnace through which an oxidizing gas flame sweeps 
with intense heat. The carbon gradually burns out until 
it is shown by test that the proper amount remains. An 
alloy of iron and manganese is added, the steel is cast into 
ingots and hammered with heavy steam hammers while 
cooling. 

The gas used in the open-hearth furnace is a sort of 
producer gas made from pure coke, air and steam. The 
following equations suggest something of its composition: 

2C + 2 = 2CO and H 2 + C = H 2 + CO. 

On account of the danger of getting impurities into the 
steel care is taken that the coke be free (or nearly so) 
from sulphur. 

As in the case of the Bessemer process there are two 
distinct methods, one with a silicon lining and the other 
with a basic lining of magnesium and calcium carbonates 
— the one for cast iron without sulphur and phosphorus 
and the other for working cast iron which contains these 
elements — so there are two open hearth methods differing 
in much the same way. 

Many high-grade steels containing nickel, chromium, 





Fig. 87. — The Converter in Action. On the left the bottom of another converter 
which has been turned over to pour out the charge of steel may be seen. By cour_ 
tesy of The Scientific American. 



BUILDING MATERIALS 



229 



manganese, and a number of the rare metals are made 
by the crucible process. These steels are used for tools, 
large guns, projectiles, etc., and are known as crucible 
steels. 

Steel contains from .8 to 2.5 per cent of carbon, while 
cast iron runs as high as 5 per cent, and wrought iron 
has practically none. Holding this middle ground in 




Fig. 88. — The Open-Hearth Furnace. The peculiar part of this furnace is the 
brick checkerwork at each end. The producer gas, which is the fuel used in this 
furnace, and air may be brought in through either checkerwork while the fire 
gases on their way to the flue pass through. the other checkerwork making the 
bricks very hot. When this condition is reached, the gas and air are brought 
in through the hot brickwork and the hot flue gases are allowed to heat the checker- 
work at the other end. By thus reversing the direction of the gases, much of 
the heat that would escape through the flues is brought back into the furnace. 

carbon content, steel also has a melting point inter- 
mediate between the other two. In other properties also 
it partakes of the good points of both wrought and cast 
iron. It can be cast, welded and hammered into shape. 
It is very strong and elastic and in addition to these 
properties it may be tempered to any degree of hard- 
ness. Tempering is accomplished by heating and sud- 
denly cooling the steel. If the steel is heated red hot 



230 



A PRACTICAL CHEMISTRY 



and then suddenly cooled, it becomes very hard. If 
cooled at a lower temperature, it is not so hard. Thus 
it will be seen that any desired hard- 
ness for any particular use may be 
obtained. 

The largest use of modern steel is 
in building the frames of buildings 
and of ships. Innumerable smaller 
uses suggest themselves to everyone. 
Compounds of Iron. — Thus far we 
have spoken of iron as a metal and 
said nothing of its compounds save 
to mention a few ores. Chemically, 
iron is not more active than many of 
the other metals, yet it forms quite 
a large number of combinations. It 
has two valences and hence forms 
89.— Hammer for two series of compounds — ferrous 
compounds in which the iron has a 
valence of two, and ferric in which 
the valence is three. Iron in the 

on the bed below. The 

hot steel is placed on the ferrous condition can be changed to 
lower half of the die. the f err j c \> y oxidation, while ferric 

The hammer with the . . 

other half fails upon it, iron by reduction is changed to fer- 

cutting the steel into the pq-itq 




Fig. 

Making Drop-forgings. 
The die is in two halves, 
one of which is placed in 
the hammer and the other 



The common acids react directly 



required shape. After 
this, the rough forging 

thus made must be w i t h j ron . hydrochloric acid, under 

trimmed and ground. . 

Many parts of machines these conditions, forms ferrous 
are thus made. chloride and sulphuric acid gives 

ferrous sulphate, but in the presence of an oxidizing 
agent the ferric salts are formed. With nitric acid iron 
forms ferric nitrate. In the air most ferrous salts change 
to ferric; a marked exception to this is the ferrous car- 
bonate, FeC0 3 , previously mentioned as the ore siderite. 
With the alkalies these salts react to form hydroxides - y 



BUILDING MATERIALS 231 

the ferrous salts forming ferrous hydroxide and the fer- 
ric, ferric hydroxide. For example, 

FeCl 2 + 2NH 4 OH = Fe(OH) 2 + 2NH 4 C1. 

(Ferrous (Ammonium (Ferrous (Ammonium 
chloride) hydroxide) hydroxide) chloride) 

FeCl 3 + 3NH 4 OH = Fe(OH) 3 + 3NH 4 C1. 

(Ferric (Ferric 

chloride) hydroxide) 

Ferrous hydroxide is a greenish-white solid which soon 
oxidizes and becomes brown in the air. Ferric hydroxide 
is a flocculent brown precipitate. 

Sulphur reacts with iron to form sulphides. The sim- 
plest of these is ferrous sulphide, FeS, which is made 
by simply heating sulphur and iron together until they 
begin to unite, after which the reaction continues with 
violence till completed. Ferrous sulphide is used to make 
hydrogen sulphide, or hydrosulphuric acid. The fer- 
rous sulphide is a heavy, dark, metal-like solid. Iron 
pyrites, FeS 2 , is the sulphide found in large quantities in 
nature and used as a source of sulphur in making sul- 
phuric acid. 

Perhaps the most used of the salts of iron is. the ferrous 
sulphate, which forms green, glasslike crystals with the 
formula, FeS0 4 .7H 2 0. It is used in the manufacture 
of ink and as a mordant in dyeing. 

"With the cyanides, iron combines to form some com- 
plicated compounds, thus the cyanide of potassium, KCN, 
unites with both ferrous and ferric cyanides to form 
potassium ferrocyanide, K 4 Fe(CN) 6 and potassium fer- 
ricyanide, K 3 Fe(CN) 6 . The first used to be called the 
yellow prussiate of potash, since it is a lemon yellow 
crystal, and the other was called the red prussiate of 



232 



A PRACTICAL CHEMISTRY 



potash. These compounds are used as reagents in testing 
for ferrous and ferric iron. With ferrous solutions, 
potassium ferricyanide forms a deep blue precipi- 
tate called Turnbull's blue, having the formula 
Fe 3 (Fe(CN) 6 ) 2 , while ferric solutions react with potas- 
sium ferrocyanide to form a deep blue precipitate of the 
formula Fe 4 (Fe(CN) 6 ) 3 , known as Prussian blue. It will 




Fig. 90. — A Lead Mine. 



be noted that the -o (ferrous) salt reacts with the -i 
cyanide and the -i (ferric) salt with the -o cyanide. 
Prussian blue is used in dyeing and for making bluing. 
Lead. — While no other metal plays so large a part as 
iron in the processes of building, several others are con- 
cerned to a considerable extent. Among these may be 
mentioned lead. This substance, more or less familiar 
to most people, is a heavy metal of a dull blue color. 
At best it does not take a high polish, and where freshly 
cut soon tarnishes in the air. It is more than eleven 
times as heavy as water. Lead is obtained from its 



BUILDING MATERIALS 



233 



rm 



EL 



most important ore, galena, or lead sulphide, PbS, by 
a very simple process. 

The ore, which is a black lustrous crystal, is heated in 
a reverberatory fur- 
nace till partly oxi- 
dized. By this oxi- 
dation the com- 
pounds lead oxide, 
PbO, and lead sul- 
phate, PbS0 4 , are 
formed. The con- 
tents of the furnace 
are then stirred to 
produce thorough 
mixing and the air 
shut off to prevent 
further oxidation of 
the mixture. The 

heating is continued and two reactions take place. The 
lead oxide reacts with the sulphide as in the equation : 



Fig. 91. — A Reverberatory Furnace. The 
fuel is burned on the grate at the left, fed by a 
draft of air which enters from below. This 
draft carries the heated gases from the fire 
over the long hearth, which is separated from 
the fire by a low wall. On the hearth is placed 
the ore or other material to be heated. Vari- 
ous types of reverberatory furnaces find many 
uses in chemical manufacturing. 



PbS + 2PbO == 

(Lead sulphide) (Lead oxide) 



3Pb _|_ SO, 
(Sulphur dioxide) 



also between the lead sulphide and lead sulphate we 
have the reaction: 

PbS + PbS0 4 = 2Pb + 2S0 2 
(Lead 
sulphate) 

Several other methods for reducing lead are also em- 
ployed. 

The softness of lead, together with its low melting 
point, 326° C, make it a useful metal for plumbing. In 
this connection it is used as pipes, sheets, and a con- 
stituent of solder, 




Fig. 92. — At, the Bottom of a Lead Blast Furnace. The slag is being drawn off 
and flows into the iron ladle. 




Fig. 93. — A Machine for Casting Pig Lead. As the large wheel turns, each mold is 
filled .with melted lead. When cold, the lead forms the pigs shown on the cars 
at the right. 



BUILDING MATERIALS 



235 




In the making of lead pipe the lead is melted and 
poured at a temperature a little above the melting point 
into a cylinder closed at the bottom by a movable piston. 
The lead is then chilled. Resting on the center of this 
piston is an iron rod of the same diameter as the bore 
of the pipe to be made. This rod passes up through a 
circular opening at the top of the cylinder. The diam- 
eter of this opening is that of the outside of the lead 
pipe. By means of a hydraulic press 
the piston is forced upward, squeezing 
the lead out through the opening 
around the iron rod, and thus a con- 
tinuous pipe is made. 

Lead, unless previously oxidized, 
withstands the action of acids to a 
marked extent, but under the influence 
of the action of the oxygen of the air 
it slowly dissolves in most acid solu- 
tions. Dilute nitric and concentrated 
sulphuric acids, being themselves oxi- 
dizing agents, dissolve lead freely. 
Owing to carbon dioxide in air, 
water containing air slowly dissolves 
lead, but impure water often produces 
a coating of lead salts on the surface 
and thus protects the rest from the action of the water. 
This question of the solubility of lead compounds in water 
is one of great importance since lead is a poison, which, 
though taken in small doses, will accumulate in the sys- 
tem and produce diseases difficult to cure. On this ac- 
count it is not wise to drink water which has been 
standing for some time in lead pipes. 

Sheet lead is used for acid tanks and chambers on 
account of its power to resist the action of acids. These 
sheets are easily made by rolling the lead out between 



Fig. 94 . — Making Lead 
Pipe. The melted 
lead in the cylinder, 
R, is forced out by the 
piston, P, through the 
opening, O, which is 
the size of the outside 
diameter of the pipe. 
The bore of the pipe 
is the size of the solid 
rod which extends up 
through O. 



236 



A PRACTICAL CHEMISTRY 



rolls. It is not ductile enough however to be made into 



wire. 



Lead enters into several alloys, one of which, solder, 
should be mentioned here. Solder is a mixture of lead 
and tin and is used to join lead or tin. 

Lead Compounds. — Lead salts of most acids are known, 




Fig. 95. — In a Lead Pipe Factory. 



though many of them are best made by the action of 
acids on the oxides of lead. These oxides are several in 
number and are commonly known as litharge, PbO, red 
lead, Pb 3 4 , lead peroxide, Pb0 2 . 

Litharge and red lead are made by oxidizing lead in 
the air. The metal is melted in furnaces through which 
air is passing and is kept constantly stirred. Litharge 
is sometimes a yellow powder and at other times a 
crystalline substance, according to the method of manu- 
facture. It is much used in glass-making, under which 



BUILDING MATERIALS 



237 



subject we shall have occasion to take it up again. In 
like manner we shall again come to red lead under 
paints. Lead peroxide can be made by oxidizing red lead 
with nitric acid. It finds an important use in storage 
batteries. 

Two other compounds of lead to be considered in the 



1 




' ' HrflH 




t 





Fig. 96. 



-Rolling Sheet Lead. The lead is carried by the small rollers and passes 
beneath the large roller which squeezes it out into a thin sheet. 



chapter on paint are lead chromate and white lead. The 
latter is a basic lead carbonate which means that it is a 
combination of lead carbonate and lead hydroxide. 
Lead acetate is called sugar of lead and is used in mak- 
ing white lead. The nitrate is another soluble salt of 
lead. It crystallizes in white crystals, and may be made 
by dissolving lead or lead oxide in dilute nitric acid. 

The sulphate of lead is a white compound almost 
insoluble in water but soluble in concentrated sulphuric 
acid. In analysis the sulphate is one of the tests for 



238 A PRACTICAL CHEMISTRY 

lead. It is obtained as a white precipitate when dilute 
sulphuric acid is added to a solution of a lead salt. An- 
other compound of lead obtained as a test for lead in 
analysis is the iodide. This is a bright yellow compound 
formed whenever a lead solution comes in contact with 
a soluble iodide. Lead iodide is quite soluble in hot 
water, from which solution it crystallizes on cooling in 
the form of iridescent scales. Lead chloride is a beauti- 
ful white compound, much more soluble in hot water 
than in cold, but of little commercial importance. Lead 
sulphide has been previously mentioned. 

Tin. — Apart from its occurrence in alloys, such as 
solder, tin is used mostly as a protective covering for 
iron. In this way it is used in large quantities to pro- 
tect sheet iron roofs, commonly called "tin roofs." In 
like manner we speak of "tin ware," meaning sheet 
iron covered with a thin coating of tin and made into 
various household utensils. 

Tinplate. — Tinplate, i. e., sheet iron covered with tin, 
is made by carefully cleaning the sheet iron and then 
dipping it into molten tin. 

Solid tin is called block tin. It is a white metal which 
is easily hammered into thin sheets called tin foil and 
used to protect many things from the action of the air. 
The student will think of numerous examples of this. In 
bars tin gives forth a crackling sound when bent. The 
melting point of tin is 235° C, nearly a hundred degrees 
lower than the melting point of lead. Tin is not so heavy 
as lead, its specific gravity being only 7.3. 

The only ore of tin is its oxide, Sn0 2 , called tin stone or 
cassiterite. This ore has been known and mined at Corn- 
wall, England, from earliest times. It is also obtained 
from a number of other countries. To obtain the metal, 
the ore must first be roasted to remove arsenic and sul- 
phur, and then heated in a reverberatory furnace with 



BUILDING MATERIALS 239 

carbon to effect reduction. The reaction may be ex- 
pressed : 

Sn0 2 + C = Sn + C0 2 
(Tin stone or (Tin) (Carbon 

Stannic oxide) dioxide) 

Tin has two valences, two and four, and hence forms 
two series of compounds. For example, its two oxides 
may be represented by the formulae SnO, stannous oxide, 
and Sn0 2 , stannic oxide. 

Not many of the compounds of tin find practical use in 
manufacturing. 

Zinc. — Zinc is another metal much used in the same way 
as tin to protect iron. 

This metal is found in many ores, such as the carbonate, 
smithsonite, ZnC0 3 ; the sulphide, zinc blende, ZnS ; the 
oxide, zincite, ZnO. It is reduced from its ores in much 
the same way as tin, except that zinc is a volatile metal 
and must be collected in a condenser. The roasted ore 
after being mixed with charcoal is put into a sort of re- 
tort with delivery tube leading into an iron receiver. The 
retort is heated and the zinc passes off as a vapor which 
condenses to a powder in the cold receiver or to a liquid 
if the receiver is hot. This reminds us of the purification 
of sulphur. 

The melting point of zinc is about one hundred degrees 
above that of lead, or at about 425° C. It boils at about 
945° C. Zinc is purified by distillation. Pure zinc is a 
white metal with a bluish luster. The physical properties 
of zinc vary greatly with the temperature. At the ordi- 
nary temperature it is often quite brittle, while at a little 
above 100° C. it becomes malleable. If rolled into sheets 
at this temperature it will retain its malleability when 
cooled. When heated above 200° C. it again becomes 
brittle. 



240 A PRACTICAL CHEMISTRY 

The largest use of zinc is in making galvanized iron. 
This is done by dipping the clean iron into molten zinc. 
The zinc adheres to the iron and crystallizes on cooling. 
Galvanized iron is much used for cornices, tanks, et cet- 
era. Zinc is also used in alloys, particularly brass, bronze, 
and antifriction metal. 

Chemically zinc is an active element, forming numer- 
ous compounds. Impure zinc reacts with most acids and 
alkalies. 

Copper. — In the history of our modern' civilization cop- 
per holds a place as prominent as that of steel, though of 
an entirely different character. In electrical construc- 
tions, i. e., in the building of dynamos and other electrical 
machines and of telephone and telegraph lines, on account 
of its ductility, malleability, toughness, and flexibility, to 
say nothing of its conductivity, no other known metal 
could well take its place. In addition to these properties, 
its resistance to atmospheric action should be mentioned ; 
when first brought into the air the bright red copper 
tarnishes with a thin coating of red-brown oxide, and 
this protects it very largely from further action. 

The reduction of copper from its' various ores is ac- 
complished in various ways, according to the character of 
the ore. Besides the native copper, which occurs in 
large quantities in the neighborhood of Lake Superior, 
the ores are mostly sulphides, oxides, and carbonates. 
The processes may be enumerated in a general way as: 
(1) the smelting of the native metal; (2) the roasting of 
carbonates and sulphides, thus forming oxides; (3) the 
reduction of oxides by heating with carbon; (4) the 
slagging off of iron and other impurities from more im- 
pure ores by fusing the ores with the proper flux. These 
processes are quite complicated and are too extensive 
to be taken up by the young chemist. In any case, the 
result is an impure copper containing slag and, usually, 



BUILDING MATERIALS 



241 



arsenic, silver, and gold. The purification of this copper 
is accomplished through the aid of the electric current. 

Electrolytic Purification. — The copper is melted and 
cast into plates. These are suspended in lead-lined tanks 
containing copper sulphate solution and are made the 
anodes, or positive electrodes. Thin plates of pure cop- 
per are suspended alternately between these as cathodes, 
or negative electrodes. As the electric current passes, 





c 


A 


G 


A 




r> 


A 




rt 


A 


C 


A 








u 


s 































































































































































































































Fia. 97. — Diagram Showing the Refining of Copper by the Electric Current. 
The heavy plates of copper marked A are the anodes of impure copper, while 
the thin sheets, C, are cathodes of pure copper. The bath is a solution of copper 
sulphate. Copper passes into solution from the anodes and is deposited upon 
the cathodes. 



the copper dissolves from the anodes and is deposited at 
the same rate on the cathodes. The impurities are not 
deposited and either fall to the bottom of the tanks or go 
into solution. The copper on the cathodes, called elec- 
trolytic copper, is melted again and cast into bars ready 
for use. It is almost pure copper. 

From the mud which collects at the bottom of the tanks 
much gold and silver are obtained. 

The uses of copper are very numerous. Besides those 
in connection with electricity we may mention its use 
in making many alloys (which will be discussed in the 
section on alloys), cooking utensils, gutters for roofs, elec- 
trotypes, and plates for engraving. 



242 A PRACTICAL CHEMISTRY 

Compounds.— vT wo oxides of copper, cuprous oxide, 
Cu 2 0, and cupric oxide, CuO, are known. The first is of 
a red brown color and is used in making red glass. The 
valence of copper in this oxide is one, while in the black 
cupric oxide it is two. Corresponding to these oxides we 
have two series of salts of copper. Thus there are two 
chlorides, the cuprous, CuCl, and the cupric, CuCl 2 . Al- 
though the copper salts of almost all acids are known, 
copper, like lead, reacts but little with acids unless pre- 
viously oxidized. This is illustrated by the use of bright 
copper in cooking utensils in which acid fruits may be 
boiled without corroding the copper so long as the steam 
from the boiling fruit protects the copper from the air. 
Copper is a poisonous metal, forming many poisonous 
compounds. Its salts are either green or blue. The 
most used of the copper salts is the sulphate, which 
forms large blue crystals of the formula, CuS0 4 .5H 2 0. 
It is much used wherever a soluble copper salt is 
wanted, as in copper plating, making insecticide solu- 
tions, et cetera. 

Aluminium. — Copper and the other metals considered 
in this section have been known from ancient times, but 
in the case of aluminium we have a gift of modern science. 
Locked in the clay and other silicates, in the cryolite and 
in the bauxite, this metal for ages remained unavailable 
till the introduction of electricity into metallurgy. With- 
in a comparatively few years this element has been 
taken out of the list of more expensive metals and has 
come into everyday use. This wonderful change has 
been brought about by the reduction process given 
below. 

We may describe Hall's process for the reduction of 
aluminium somewhat as follows. An iron box (see Fig. 
98) is lined with carbon and joined to the negative wire, 
making it the cathode, or negative electrode. Into this 



BUILDING MATERIALS 



243 



the mineral cryolite, which is a compound of sodium, 
aluminium, and fluorine of the formula, Na 3 AlF 6 , is put in 
lumps and packed around a number of carbon anodes 
suspended on copper rods. The current is turned on and 
enough heat is generated to melt the cryolite. The alu- 
minium ore, which is the oxide (A1 2 3 ) made from bauxite 
(H 4 A1 2 5 ), is next added. This dissolves in the cryolite, 
and, through the action of the electric current, is separ= 




Fig. 98. — Hall's Reduction Furnace for Aluminium. C is the carbon lining 
which serves as the cathode while the carbon rods, A, A, are anodes. The liquid 
is molten cryolite in which the aluminium oxide has been dissolved. The molten 
aluminium goes to the bottom and is drawn off into molds at T. 



ated into aluminium, which collects at the bottom of 
the box, and oxygen that goes to the anodes and, re- 
acting with the carbon, forms carbon monoxide. The 
molten metal is drawn off into molds and hardens into 
bars. 

Aluminium is a white metal with a characteristic bluish 
tint. It is much lighter than iron, having a specific grav- 
ity of only 2.6. It may be rolled into sheets and foil, 
drawn into wire, and otherwise worked as other metals. 
Its melting point is 700° C. 

Bright aluminium scarcely seems to tarnish in the air, 
hut careful examination shows it to be covered with a 
thin coating of oxide which protects it from the air. 
IVlien powdered and mixed with an oxidizing agent the 
metal may be burned with great violence and a very high 
temperature produced. Advantage is taken of this fact 



244 



A PRACTICAL CHEMISTRY 



in thermit, which is a mixture of powdered aluminium 
and oxide of iron used to produce molten iron quickly for 

mending castings, weld- 
ing street car rails, et 
cetera. The thermit is 
put into a crucible with 
a bottom that may be 
quickly removed. The 
crucible is placed over 
the pieces of iron to be 
welded. By the aid of 
some magnesium pow- 
der or other combusti- 
ble material the mix- 
ture in the crucible is 
lighted. In a few min- 
utes the aluminium 
has combined with the 
oxygen of the iron 
oxide, forming alumin- 
ium oxide, A1 2 3 , leav- 
ing the iron free and 
in a molten condition. 
The bottom of the cru- 
cible is then opened 
and the molten iron 
permitted to run over 




Fig. 99. — Use of Thermit in Welding the 
Stern Shoe of a Steamship. In this pro- 
cess, the broken parts are placed in their 
proper positions and surrounded by a mold. 
A crucible, the bottom of which can be 
suddenly opened, is filled with thermit and 
supported above the break. The thermit 
is lighted and soon produces a quantity of 
molten iron. The crucible bottom is then 
opened, the molten iron pours into the mold 
and the broken parts are welded together. 
The repair of the stern shoe shown in the 
figure was accomplished in less than three 
days — a saving of time which meant the 
saving of much money to the owner. 



the parts to be welded, 
thus: 



The reaction may be expressed 



2A1 + Fe 2 3 = A1 2 3 + 2Fe 



Powdered aluminium is used in the same way to reduce 
other metals from their oxides. This use of aluminium as 
a reducing agent has become quite extensive and is called 
the Goldschmidt process. 




Fig. 100. — Carborundum Furnace. In the manufacture ol carborundum at Niagara 
Falls electric furnaces are employed. The outside of one of these furnaces is 
shown in the picture and a cross-section of the same in Figure 101. The furnace 
consists of a rectangular brick structure into which is put a thick layer of the 
materials to be heated. These include sand, coke, sawdust and common salt. 
Upon this layer a core of loosely placed coke is laid between the carbon elec- 
trodes. More of the material is put on top of this. The electric current is 
passed through this mass for a number of hours. The intense heat thus produced 
brings about a number of chemical changes, the most important of which may 
be represented by the equation: 

Si02 + 3C = SiC + 2CO 
silicon dioxide + carbon = carborundum + carbon monoxide. 
It will be observed that the silicon of the sand is reduced by carbon and at 
the same time combines with carbon forming carborundum, which is chemically 
known as silicon carbide. Carborundum, as it comes from the furnace, is often 
in the form of beautiful crystals of many colors. It is next to the diamond in 
hardness and is used for grinding and polishing in all conceivable ways. So 
numerous are the uses to which it is applied and so large is the demand for this 
substance that many million pounds of it are produced at Niagara Falls each year # 
In Figure 102 we have a collection of carborundum products. 



-246 



A PRACTICAL CHEMISTRY 




Fig. 101. — Cross-section of Carborundum 
Furnace. 



On account of its lightness, great strength, and resist- 
ance to atmospheric action, together with the many ways 
in which it may be worked, aluminium finds a great va- 
riety of industrial uses. We see it serving as a decoration 
on the outside of books, the material of household uten- 
sils, a decorative and protective covering of steam pipes, 
radiators, and mail boxes, as automobile castings, and 

wire for conduct- 
ing electricity, to 
say nothing of a 
thousand smaller 
uses. 

Compounds. — 
Among the com- 
pounds of alumin- 
ium, its oxide 
stands first in simplicity and perhaps first in impor- 
tance, since from it the metal is obtained. The mineral 
corundum, which is used as emery, is aluminium oxide. 
Emery has been much used for grinding and polishing in 
the form of emery wheels, cloth, and powder. Carborun- 
dum is a manufactured substitute often used at present 
in place of emery. When pure, aluminium oxide is white, 
but it may be easily colored by small quantities of other 
metals. The sapphire and ruby, the oriental amethyst 
and emerald are minerals containing aluminium oxide 
thus colored. 

Hydroxide of Aluminium. — The hydroxide of alumin- 
ium, Al(OH) 3 , may be made by the action of alkaline 
hydroxides as ammonium hydroxide upon the aluminium 
salts. It is a white, jelly-like substance insoluble in 
water but soluble in sodium and potassium hydroxide 
and in acids. Aluminium hydroxide may also be made 
by the action of alkali carbonates on aluminium salts. 
This takes us back to baking powders made of sodium 



BUILDING MATERIALS 247 

bicarbonate and alum, and we find the explanation of 
their action in the fact that aluminium hydroxide and 
carbon dioxide are formed thus: 

NaAl(S0 4 ) 2 + 3HNaC0 3 = 3C0 2 + 2Na 2 S0 4 + 
(Alum) (Acid sodium (Sodium 

carbonate) sulphate) 

Al(OH),. 

(Aluminium 
hydroxide) 

With most of the acids aluminium forms salts, though 
it does not dissolve easily when pure in nitric or sulphuric 
acids. It is quite soluble in sodium and potassium hy- 
droxides, hence alkalies and soaps containing free alkali 
should not be used in cleaning aluminium ware. Among 




Fig. 102. — Carborundum Products. 

the salts of this metal the alums have previously been 
mentioned and explained. 

Both alum and aluminium sulphate are much used in 
dyeing and purifying water. 

Aluminium Chloride. — Aluminium chloride, A1C1 3 , is 
made by the direct action of the metal on hydrochloric 
acid. It is used as a reagent in organic chemistry. 



248 A PRACTICAL CHEMISTRY 

One silicate of aluminium is of much importance and 
will be considered in another chapter. 

SUMMARY 

The building materials discussed in this chapter may be included 
under the headings, Materials Required by the Mason and 
Metals Used in Building. Under the former are included 
building" stones, bricks, mortar, and cement. 

Building stones are composed of silicates, silicon dioxide and 
carbonates. The silicates are quite complicated and are 
the salts of various silicic acids. The metals in these 
silicates are mostly aluminium, calcium, magnesium, sodium, 
potassium, and iron. The carbonate rocks are mostly cal- 
cium and magnesium carbonates. 

Bricks are made from clay, which is mostly an aluminium silicate 
mixed with sand and small quantities of other compounds. 

Mortar is a mixture of slaked lime, sand, and water. The slaked 
lime is made from limestone (calcium carbonate) by two 
operations; first the carbonate is decomposed by heat, form- 
ing quicklime, or calcium oxide and carbon dioxide; then by 
adding water to the quicklime slaked lime, or calcium hydrox- 
ide, is formed. Sand is silicon dioxide. Mortar hardens by 
the absorption of carbon dioxide and the evaporation of 
water. 

Cement is made by heating together a mixture of cement rock and 
limestone and grinding the clinker thus formed into a fine 
powder. Cement hardens on account of a chemical change 
produced by the water. 

Plaster is a form of mortar applied to the surface of walls. 

Plaster of Paris is calcium sulphate from which part of the 
water of crystallization has been removed by heating. It 
hardens by taking up water and recrystallizing. 

The metals mentioned include iron, lead, tin, zinc, copper, and 
aluminium. 

Iron is reduced from its ores, which are mostly oxides or car- 
bonates, by means of the blast furnace. The product of 
this furnace is pig iron, which is a form of cast iron. It 



BUILDING MATERIALS 249 

contains, besides iron, silicon, carbon, sulphur, phosphorus, 
and small quantities of a number of other elements. 

Wrought iron is the purest iron used for commercial purposes. 
It has a high melting point and is the product of the pud- 
dling furnace. Most steel is made by one of two processes, 
the Bessemer and the open hearth. 

High-grade steels containing various other metals are made by 
the crucible process. Steel has many of the good properties 
of both cast and wrought iron, and may also be tempered 
to any degree of hardness. Iron forms both ferrous and 
ferric compounds, corresponding to the two valences, two 
and three. These include ferrous and ferric salts of nearly 
all acids. 

Lead is obtained from the sulphide, PbS, by the air reduction 
process. Lead is used for making pipe, sheet lead, and 
alloys. It forms several oxides and many useful salts. 

The important ore of tin is its oxide, Sn0 2 , from which it is 
obtained by reduction by carbon. It is used as tinplate, 
which is iron coated with tin, tinfoil, block tin, and in alloys. 
Tin forms stannous compounds, in which its valence is two, 
and stannic, in which its valence is four. A few of these 
are of practical importance. 

Zinc is largely used as galvanized iron, made by dipping clean 
iron into molten zinc. It is also used in many alloys. 

Impure copper is obtained from its ores by complicated proc- 
esses. It is purified by electrolysis. The properties of cop- 
per make it particularly useful in electrical constructions. 
Like lead, copper is not easily acted upon by most acids. 
Its salts are poisonous. 

Aluminium is a constituent of many minerals. It is obtained 
from bauxite by Hall's process. It finds many uses as a 
metal because of its lightness, great strength, and resistance 
to atmospheric action. In powder form it is used as a re- 
ducing agent in the Goldschmidt process. 

REVIEW QUESTIONS 

1. How does silicon resemble carbon? 

2. Why does silicon dioxide have so many names? 



250 A PRACTICAL CHEMISTRY 

3. Explain the composition of the silicate rocks. 

4. How do you explain the formation of the chalk deposits 
of England? 

5. How does coral differ from calcite? 

6. How can you make lime water from limestone? 

7. Name the four most important elements in a brick 
wall. 

8. What would you obtain by heating a plaster cast? 
- 9. How would you tell whether shot is a metal? 

10. Mention something which could not be built of concrete, 
and give your reason for thinking so. 

11. What is gained by having blast furnaces built so high? 

12. What compound will be found in the slag when the flux 
is limestone and the ore contains silicon dioxide? 

13. What is the reducing agent in the blast furnace? 

14. How could you tell whether a given sample is cast iron, 
wrought iron, or steel? 

15. How would Bessemer steel without the spiegeleisen differ 
from wrought iron? 

16. How could you remove the temper of hard steel? 

17. How could you make ferric hydroxide from siderite? 

18. How can salts of lead be made? 

19. Why is the solubility of lead compounds an important 
matter? 

20. Why cannot iron pipe be made in the same way as lead 
pipe? 

21. Upon what property of tin does its most important use 
depend ? 

22. How does the purifying of zinc differ from the purifying 
of other metals? 

23. How could you tell whether a soluble compound contained 
lead or zinc? 

24. How would you make electrolytic copper from copper 
carbonate ? 

25. How are copper salts used by fruit growers? 

26. Why is aluminium used in paints? 

27. How could you use powdered aluminium to reduce such 
metals as chromium? 






BUILDING MATERIALS 251 

28. Why does aluminium make good cooking utensils'? 

29. When alum is used in baking powder what compounds 
does it leave in the bread? 

30. How could you make artificial emeralds? 



CHAPTER XX 

ALLOYS 

Not satisfied with the numerous metals provided for 
him, man has obtained a large number of substances by 
melting together mixtures of two or more metals. The 
substances thus formed are called alloys. The metals melt 
and dissolve in one another so that the alloy is not a mere 
mechanical mixture, but is governed by the laws of solu- 
tion. Part of the metals also sometimes form chemical 
compounds with one another, and these compounds dis- 
solve in the remaining metal. On this account it is not 
possible to make alloys of all metals in all proportions 
as we may make mechanical mixtures in all proportions. 
Some of these alloys are of great commercial value, and 
find uses which could scarcely be filled by the metals 
themselves. Take as an illustration brass, which is com- 
posed of about two-thirds copper and one-third zinc. 
Brass has many desirable qualities which fit it for a great 
variety of uses. It is hard, malleable, and ductile, and 
can be cast. It is capable of taking a high polish and 
has a beautiful goldlike color. It is quite elastic. Its 
most objectionable property is its odor. The composition 
of brass, within moderate limits, is quite variable and a 
number of different names such as Dutch metal, Muntz 
metal, and pinchbeck are applied to the different kinds. 
"When left exposed to moist air, brass becomes coated 
with a green deposit of basic copper carbonate. This is 
sometimes spoken of as verdigris, though the true verdi- 
gris is an acetate of copper. 

252 



ALLOYS 253 

Closely related to this alloy is another called bronze. 
This again, like brass, is of variable composition so that 
quite a number of different bronzes are made to suit 
various purposes. In all cases the main constituent is 
copper, while tin and zinc are usually the others. Some- 
times small amounts of lead and iron are used. Traces 
of phosphorus are found in phosphor bronze, and of sili- 
con in silicon bronze. 

Bronzes are usually of a reddish brown color; their 
beauty and durability make them particularly desirable 
for statues, tablets, medals, and coins. 

Aluminium Bronze. — Aluminium bronze is a modern 
alloy of considerable importance. Its composition, un- 
like the other bronzes, is copper and aluminium, the lat- 
ter constituting about 10 per cent, of the alloy. It has 
a yellow color, is elastic, and resists the action of the 
air. It has been used in making hulls for yachts. 

Bell Metal. — Bell metal, the material of which bells are 
made, is usually a bronze of copper and tin of varying 
proportions. The tin is usually from twenty to twenty- 
five per cent of the alloy. 

White Metal. — Several different sorts of white alloys 
are used for the foundation metal in plated spoons, forks, 
et cetera. These metals are largely tin with a little cop- 
per and sometimes antimony. A similar alloy, Britannia 
metal, was in former years much used for table ware. 
It contains about 90 per cent tin, 2 per cent copper, and 
8 per cent antimony. Another old alloy used for the 
same purpose was pewter. It is composed of tin and 
lead. Solder is of the same composition, the amount of 
tin varying from about one-third to one-half. 

Speculum Metal. — Speculum metal is an alloy richer 
in tin and might almost be classed as a bronze. It is 
used to make the reflectors of telescopes. 

An alloy much used in scientific instruments is Ger- 



254 



A PRACTICAL CHEMISTRY 



man silver. It does not contain any silver but is about 
one half copper and the other half nickel and zinc. It 
looks somewhat like silver. 

Magnalium. — Magnalium is an alloy of aluminium and 
magnesium with from 75 to 90 per cent of the former. 
It is harder than aluminium but not so heavy. 

Another class of alloys quite complicated as to com- 
position is termed antifriction metals, i. e., metals used 
to put into the boxes of machinery to prevent friction. 
Babbitt's metal is a well-known member of this class. 




.. 



<P) 



O 



Fig. 103A. Fig. 103B. 

Fig. 103A. — Fusible Links. The two sections A and B of the link are soldered to- 
gether with an alloy which melts at a very low temperature. This link is made a 
part of a chain or other fastener which prevents a heavy fire-proof door from 
swinging shut. When a fire occurs, the heat melts the alloy and the two parts 
of the link separate as shown in Figure 103B, thus allowing the door to shut. 

A very interesting class of substances . is known as 
fusible metals. These are made of various proportions 
of lead, antimony, cadmium, tin, and bismuth. They may 
be made to melt at any desired temperature, even far 
below the boiling point of water. Many important uses 
are found for such alloys. For example, the fuses in 
electrical circuits which melt when too great a current 
is passing, and the plugs in automatic fire extinguishers. 
Likewise the fastenings which hold open fireproof 
doors between rooms in large buildings contain sections 
of fusible metal. By the heat of the fire these fastenings 
are melted and the doors swing shut. (See Fig. 103.) 

Many of these alloys also have the property of expand- 



ALLOYS 



255 



ing before they harden, and so may be used to make 
casts of dies. This property of expanding during cool- 
ing is derived from the antimony which they contain. 
Type metal, which is an alloy of lead, tin, and antimony, 
gets its value from this property. 

Many other alloys are in everyday use. Among these 




Fig. 104. — Linotype Machine in Which Type Metal is Cast into Type by Press- 
ing the Letters on the Keyboard. 



may be mentioned shot, which is an alloy of lead with 
a little arsenic. The coins of all nations are alloys. 
The United States gold is 90 parts gold and 10 
parts copper, while our silver coins are also 10 per cent 
copper. 

Amalgams. — Amalgams are alloys containing mercury. 




Fig. 105. — Coining Money in the United States Mint, Philadelphia. 



ALLOYS 257 

SUMMARY 

Alloys. — Alloys are substances formed by melting together two or 

more metals. They often have desirable properties not 

possessed by single metals. Among alloys having color in- 
clude : 

Brass, containing copper and zinc. 

Bronze, containing copper, tin, zinc, and sometimes lead and 
iron. 

Phosphor bronze, bronze containing a little phosphorus. 

Silicon bronze, bronze containing a little silicon. 

Aluminium bronze, containing copper and aluminium. 

Bell metal, containing copper and tin. 
White Metals. — White metals include: 

Britannia, composed of tin, copper, and antimony. 

Pewter, containing tin and lead. 

Solder, containing tin and lead. 

Speculum, containing copper and tin. 

German silver, containing copper, nickel, and zinc. 

Magnalium, containing aluminium and magnesium. 
Antifriction Metals. — Antifriction metals may be illustrated by 

the well-known Babbitt's metal. 
Fusible Metals. — Fusible metals contain lead, antimony, cad- 
mium, tin, and bismuth. 
Shot. — Shot contains lead and arsenic. 
United States Gold Coin. — United States gold coin contains gold, 

90 per cent, and copper, 10 per cent. 
United States Silver Coin. — United States silver coin contains 

silver, 90 per cent, and copper, 10 per cent. 
Amalgams. — Amalgams contain mercury and other metals. 

REVIEW QUESTIONS 

1. Is it an easy matter to devise a new alloy? 

2. Brass often has nearly the same color as gold; why is it 
not equally desirable? 

3. How could you tell aluminium bronze from brass? 

4. What alloy has a composition similar to bell metal? 



258 A PRACTICAL CHEMISTRY 

5. Under what conditions would there be danger of poisoning 
from pewter dishes? 

6. What properties should a good solder have? 

7. How is a reflecting telescope constructed? 

8. How would you use Babbitt's metal on old machinery? 

9. State where you have observed fusible metal fastenings 
holding open fireproof doors in your school building or in other 
public buildings. 



CHAPTER XXI 
METALS USED FOR PLATING AND DECORATING 

Only a few metals are included under this class, namely, 
gold, platinum, silver,- nickel, and copper. 

Gold. — From the earliest times men have given every- 
thing and risked even life itself for this metal. This is 
not so much due to its scarcity, for other metals are found 
in even smaller quantities, as it is to its beauty, its re- 
sistance to chemical action, and its other desirable quali- 
ties. 

Gold is found in many parts of the earth in small quan- 
tities, mostly in the native condition, i. e., not combined 
with other elements. It is, however, often alloyed with 
silver, and is sometimes found associated with copper and 
also with the rare element tellurium. In a few places 
large nuggets are found, but more often it occurs as gold 
dust ; that is, in very fine particles. These particles are 
sometimes mixed with sand and gravel, while in other 
places they are scattered through quartz rock. 

Gold is mined in many different ways. Wherever the 
gold is simply mixed with sand and other earthy materials 
various kinds of hydraulic mining are employed. On a 
large scale this is done by forcing streams of water 
against the mountainside, washing the earth away 
through long sluices in which are little pockets of mer- 
cury. The gold, being more than 19 times as heavy as 
water, soon sinks into the pockets where it dissolves in 

259 



METALS USED FOR PLATING AND DECORATING 261 

the mercury, forming what is known as an amalgam — 
all alloys of mercury are called amalgams. The mer- 
cury, being volatile, may be removed from the gold 
amalgam by distillation. Gold-bearing quartz is treated 
by the cyanide process. The rock is crushed to a powder 
and then mixed with a dilute solution of sodium or 
potassium cyanide. The cyanide reacts chemically with 
the gold, forming a double cyanide of gold and potassium 
(or sodium, if sodium cyanide is used), which may be 
called potassium aurocyanide. The reaction requires oxy- 
gen from the air or an oxidizing agent and may be ex- 
pressed : 

2Au + 4KCN + + H 2 = 2KAu(CN) 2 + 2KOH. 
(Potassium (Potassium 

cyanide) aurocyanide) 

The gold is removed from this solution by electrolysis, 
i. e., the electrodes of an electric circuit are put into the 
solution, and the current started. The gold collects on 
the cathode. 

The other methods of mining gold need not be men- 
tioned except in the case of those ores which require 
smelting. These include gold-bearing copper ores, which 
give copper containing gold. The separation of this gold 
and copper has previously been mentioned in the electro- 
lytic refining of copper. 

Gold is a heavy yellow metal, very malleable and duc- 
tile. It may be drawn into the finest wire and hammered 
into the thinnest foil. Its uses are innumerable. The 
most important is for coin. Enormous quantities are also 
used in jewelry, but always as an alloy containing much 
copper or silver or both. The proportion of gold, how- 
ever, in jewelry is not expressed in per cent, but in 
carats, or twenty-fourths. For example, if a ring is 
marked 18 K. it is intended to indicate that it is eighteen 



262 A PRACTICAL CHEMISTRY 

twenty-fourths, or three-fourths gold. Much jewelry is 
only 14 carat gold, and some even so low as 10 carat. 
Gold is a very soft metal, and watches and other articles 
made of 18 carat gold do not wear well. 

Dentistry requires much gold each year. A consider- 
able amount is also used as gold leaf for decorating, 
though much decorating is done with foil made of Dutch 
metal. 

Compounds. — Gold does not form many compounds. 
The chloride, AuCl 3 , is much used in photography. It is 
made by dissolving gold in aqua regia, which is a mixture 
of hydrochloric and nitric acids. 

Gold Plating*. — Potassium aurocyanide has been men- 
tioned in connection with the cyanide process. It is also 
used for gold plating. The object to be plated is made 
the cathode, while the anode is made of pure gold. Thes^ 
are put into a bath of potassium aurocyanide and the cur- 
rent turned on. After the plating is deposited it must be 
polished. 

Platinum. — Platinum is a heavy tin colored metal which 
is found mixed with other metals, such as the rare ele- 
ments osmium, iridium, and palladium. On account of 
its high melting point and its resistance to chemical ac- 
tion, together with the fact that it expands at about the 
same rate as glass, platinum is much used in making 
scientific apparatus. For example, in chemistry it is used 
for crucibles and dishes and wherever it is necessary to 
fuse a wire into glass. In electric light bulbs we find 
bits of platinum wire penetrating the glass to carry the 
current. These numerous uses, together with its limited 
supply, have brought the price of platinum above that of 
gold. 

In addition to the uses just mentioned, it is found 
valuable in an entirely different line, i. e., as a catalytic 
agent, as we have previously studied under the manu- 






METALS USED FOR PLATING AND DECORATING 263 

facture of sulphuric acid. For this purpose finely di- 
vided platinum is used, which is obtained by the decom- 
position of platinum compounds. The metal dissolves in 
aqua regia, forming chlorplatinic acid, H 2 PtCl 6 . This 
forms salts with the alkalies as (NH 4 ) 2 PtCl 6 , ammonium 
chlorplatinate. If this salt is heated to redness it de- 
composes and spongy platinum is left, which is one form 
of finely divided platinum. Another form of much more 
finely divided platinum is called platinum black. It is 
made by the reduction of chlorplatinic acid, and is also 
used as a catalytic agent. 

Silver. — Nearly all galena contains silver sulphide. Sil- 
ver sulphide also occurs alone, and is called silver glance. 
These, together with the copper ores containing silver, 
are the most important sources of this metal. 

Silver is reduced from the sulphide in the same way 
that lead is obtained from galena, hence lead usually con- 
tains silver when first reduced from the ore. The silver 
is usually removed from the lead by Parkes' process, 
which depends upon the fact that silver alloys more 
readily with zinc than with lead. The lead is melted and 
run out into large pans, and the proper amount of zinc is 
added. This collects a considerable amount of the silver 
and the alloy thus formed, being lighter than lead, floats 
to the top. Again this alloy melts at a higher tempera- 
ture than lead, and hence begins to solidify before the 
lead. It is removed by means of skimmers with as little 
of the lead as possible. The alloy is squeezed in a hy- 
draulic press to remove the lead, and then put into a re- 
tort and the zinc distilled. This leaves the silver con- 
taminated with lead and containing a little gold, since 
gold is often in the lead. To remove the lead from the 
silver the mixture is cupelled; that is, it is heated with 
a current of air in a reverberatory furnace, the hearth of 
which is lined with bone ash. The lead oxidizes and is 



264 



A PRACTICAL CHEMISTRY 



absorbed by the bone ash, passes off as vapor, or runs 
away as molten lead oxide. An alloy of silver and gold 
remains. These are separated by electrolysis. The al- 
loy is made into anodes, these are surrounded with silken 
bags and put into a bath of silver nitrate solution. Pure 
silver cathodes are put between the anodes. When the 




Fig. 107. — Desilvering Lead. The alloy^of zinc and silver is being collected from 
the surface of the molten lead. 



current is turned on the silver dissolves from the anodes 
and crystallizes in beautiful white crystals on the 
cathodes. The gold is caught as a black powder in the 
silken bags. The silver crystals are washed and melted 
into ingots. The gold is treated in like manner. 

The purity of silver and gold is expressed in parts per 
thousand, and is said to be so many "fine"; thus the 
United States silver coin is 900 fine, i. e., 900 parts silver 
to 100 parts copper. Sterling silver is 925 fine. 

Silver is used for coins, table ware, jewelry, and in 
photography. Much table ware is plated. Silver plating is 



METALS USED FOR PLATING AND DECORATING 265 

done in much the same way as gold plating. The solution 
used is potassium argenticyanide, KAg(CN) 2 . Pure sil- 
ver plates are used as anodes, and the ware to be plated 
is made the cathodes. The ware must be frequently pol- 
ished during the process. 

Silver is a white metal capable of being polished and 
yet is not too hard to be easily worked into any desired 
c 

J 




Fig. 108. — Silver Plating. The articles to be plated are hung as cathodes from 
the negative wire, while between them pure silver anodes are suspended on the 
positive wire. The bath is a solution of argentic potassium cyanide. 

form. It melts at about 954° C. It is an active element 
chemically and forms salts with many acids. Most of its 
salts are insoluble. The nitrate, however, which is made 
by dissolving silver in nitric acid, is quite soluble and 
is much used where a soluble silver compound is wanted. 
Silver is easily acted upon by sulphur, which turns it 
black. The small amount of sulphur compounds in the 
air soon blackens silverware, while spoons used in eggu 
often are tarnished by the sulphur of the eggs. 

Silver in photography will be considered in another 
chapter. 

Nickel. — This metal is found in comparatively few 
places, and the methods used to extract it from its ores 
are quite complicated. On account of its beautiful white 
color, the high polish which it can take, and its resistance 
to the action of the air, nickel is much used for plating, 



266 A PRACTICAL CHEMISTRY 

particularly water pipes, faucets, and other plumbing 
supplies. It is best when plating iron to coat it first with 
copper. For nickel plating a bath of nickel ammonium 
sulphate is used, while a nickel plate is made the anode 
and the object to be plated the cathode. 

Nickel is also used as an ingredient of coins and other 
alloys, and of nickel steel. 

Copper Plating. — Copper plating is done in much the 
same way as nickel plating, but is of much wider applica- 
tion. The bath is a solution of copper sulphate, the anode 
is pure copper, and the object to be plated is the cathode. 

Electrotyping. — Electrotyping is one application of cop- 
per plating. An impression of the type is made in wax, 
which is coated over with graphite to make it a con- 
ductor of electricity. The wax is then made the cathode 
as in other copper plating and coated with copper. When 
removed from the bath the wax is melted out, and a cop- 
per shell the same shape as the type remains. The elec- 
trotype is next reinforced by filling in the back with an 
easily melted alloy. 

SUMMARY 

Gold. — Gold is highly valued because of its beauty, its resistance 
to chemical action, and its scarcity. It is mined by the 
hydraulic process, which consists in washing away the earth 
and catching the gold in pockets of mercury, and by the 
cyanide process, in which the quartz is crushed to powder 
and then treated with a dilute solution of sodium or potas- 
sium cyanide. Gold is a soft metal, and when used is alloyed 
with copper or silver. The proportion of gold in jewelry is 
expressed in carats, or twenty-fourths. 

A mixture of hydrochloric and nitric acids is known as 
aqua regia. It will dissolve gold and form gold chloride, 
AuCl 3 . 

Platinum. — Platinum is a very heavy, tin colored metal, highly 
valued because it resists chemical action. In a finely divided 
condition it is a catalytic agent* 



METALS USED FOR PLATING AND DECORATING 267 

Silver. — Silver often occurs as the sulphide along- with lead sul- 
phide, and is reduced along with lead. The lead is then 
desilvered by the Parkes' process. The purity of silver and 
gold is expressed in parts per thousand, or "fine." Silver 
becomes blackened by the sulphur compounds in the air. 
Nickel is a white metal, not easily acted upon by the air. 
It is much used in plating faucets, et cetera. 

Copper. — Copper is used in plating iron which is afterward to 
be plated with nickel, and for electrotyping. 

In plating with these various metals a plate of the pure 
metal is made the anode, while the object to be plated is 
made the cathode. When plating with copper the bath is a 
solution of copper sulphate; with nickel the solution is 
nickel ammonium sulphate; while in the case of platinum, 
gold, and silver the cyanides are used dissolved in a solution 
of potassium cyanide. 

REVIEW QUESTIONS 

1. How can gold be separated from copper? 

2. How would you get gold out of a sand pile? 

3. How is the manufacture of sodium cyanide connected 
with the electrolysis of salt? 

4. With gold at $20 per ounce, what is the value of an ounce 
of 10K gold? 

5. Which properties of platinum make it useful in the 
electric light bulb? 

6. In what two ways is platinum used in sulphuric acid 
manufacture ? 

7. How could you remove silver from an alloy of zinc and 
silver? 

8. How could you remove lead from an alloy containing 
lead and gold? 

9. Why do silver mustard spoons turn dark when in use? 

10. How does plating with nickel differ from plating with 
silver? 

11. How does electrotyping differ from ordinary copper plat- 
ing? 



CHAPTER XXII 
GLASS AND POTTERY 

At a date so early that history fails to record it man 
discovered glass, and even in ancient times fine varieties 
of glass were made. Science, however, failed to come to 
the aid of this industry to replace its crude and ignorant 
methods by intelligent ones until quite recently, and 
even now it is not in this country but in Europe that 
great progress has been made. 

A great variety of substances pass under the name 
of glass. In general, we may say that glass is a mixture 
of silicates more or less transparent. In the chapter 
on silicates we learned of water glass, and how it is 
made by melting together sand and carbonates of sodium 
and potassium. It was further stated that this substance 
is called water glass since it is glasslike and soluble 
in water. Now it happens that all glass contains 
either sodium or potassium silicate mixed with a suffi- 
cient quantity of insoluble silicates to render it insol- 
uble, calcium and lead silicates being the ones used for 
this purpose. 

In making glass we follow a plan similar to that for 
making water glass. Into a large fire clay pot the batch 
(or mixture of substances from which the glass is to be 
made) is put. This mixture may contain the following 
substances : sand, sodium carbonate, potassium carbonate, 

268 



GLASS AND POTTERY 269 

limestone, sodium nitrate, sodium sulphate, lead oxide, 
and small quantities of numerous other metallic oxides, 
though all of these substances are not used in the same 
batch. 

In making window glass the mixture would be sand, 
limestone, sodium and potassium carbonates, sodium ni- 
trate, and perhaps a very little manganese dioxide. The 
nitrate is put in as an oxidizing agent to remove impuri- 
ties. The fire clay pot is gradually heated till the mix- 
ture begins to melt, the temperature is then kept nearly 
constant for some time to let the gases escape quietly. 




Fig. 109. — Exhibit of the Process of Making Glass. The row of bottles is in- 
tended to show the various ingredients in their relative quantities required for 
making glass — the smallest bottle representing the required quantity of man- 
ganese dioxide, while the one marked "batch" is the mixture of all. "Cullet" 
is broken glass. It is added to serve as a flux which will hasten the melting of 
the other materials. 



Later the temperature is increased till the glass becomes 
a thin liquid. The impurities which have come to the 
top are removed as a scum. The glass is next permitted 
to cool until it takes on a pasty condition. In this con- 
dition it is ready to be worked. Chemically the manu- 
facture of glass consists in uniting silicon dioxide with 
metallic oxides, thus forming silicates, while the sub- 
stances previously united with these metallic oxides pass 
off as gases or vapors. Thus : 



270 



A PRACTICAL CHEMISTRY 



CaC0 3 + 
(Calcium 
carbonate) 


Si0 2 
(Silicon 
dioxide) 


CaSi0 3 
(Calcium 
silicate) 


+ co 2 

(Carbon 
dioxide) 


and 








Na 2 SO, + 
(Sodium 
sulphate) 


Si0 2 = 


Na 2 SiO s 
(Sodium 
silicate) 


+ so 3 . 

(Sulphur 
trioxide) 



Glass is worked into articles of an infinite variety of 
shapes by blowing and molding. When the glass is in 
a pasty condition a lump of it is taken upon the end of 
a blowpipe and air is forced into it either by the mouth or 
by machinery. When window glass is blown by mouth 
the glass forms a hollow globe, and when this is swung 
back and forth like a pendulum it becomes a cylinder. 
These cylinders are next cut from end to end and the 
ends removed. They are then placed upon a heated sui> 
face, where they flatten out into large panes which are 
then trimmed to the required size. In making bottles, 
a mold made in two halves hinged together is used. 
The mold is closed on a lump of glass on the end of a 
blowpipe. When air is forced through the pipe the glass 
fills the mold and takes its shape. Any letters or designs 
which may be upon the mold appear on the bottles. 

Many small objects, such as dishes and pitchers, are 
made of pressed glass ; that is, the glass is molded without 
blowing. 

The large, thick plate glass used for show-windows is 
made by pouring out the molten glass on a table and 
rolling it to the desired thickness with a roller. It is then 
polished. 

Almost all glassware, particularly those pieces which 
are thick and of irregular shape, if cooled quickly con- 
tract unequally while cooling, and hence are under a 
strain from which they are apt later to break. To avoid 



GLASS AND POTTERY 271 

this the glass must be annealed. This is done by passing 
it slowly through a long oven, which is quite hot at one 
end but becomes cooler and cooler until at the other end 
it is no warmer than the outside air. Glass that has not 
been annealed may fly to bits on receiving a mere scratch. 
This is well illustrated by the "Rupert's drops" which 
are drops of glass that have been quickly cooled. If 
these are scratched they break to pieces with explosive 
violence. 

The various properties of the different kinds of glass 
depend largely upon its composition. Glass containing 
much sodium melts at a lower temperature than that 
containing potassium. Also that with much alkali is 
more affected by water and steam. Again, by substitut- 
ing lead oxide for limestone and potassium carbonate for 
sodium carbonate in the batch, the very brilliant "flint" 
glass is obtained. This is the glass used for cut glass 
dishes. 

In most cases the colors of glass are obtained through 
the addition of metallic oxides. Blue glass is made by 
adding cobalt oxide, chromium compounds give green, 
while ruby glass is obtained from the use of metallic gold 
or cuprous oxide. AVhen cuprous oxide is used, care is 
required lest any oxidizing substances be present, since 
these would oxidize the copper oxide to cupric oxide and 
this would produce an entirely different result. Iron gives a 
green tint which may be counteracted by manganese. 

Pottery. — In a previous chapter we have learned of 
clay as aluminium silicate derived from the weathering 
of feldspars and the washing away of the alkali silicates. 
The purest clay, then, would be composed entirely of 
aluminium silicate, to which the formula Al 2 Si 2 7 .2H 2 
is ascribed. This is called kaolin and is a white powder. 
Feldspars are seldom pure and all of their insoluble im- 
purities remain in the clay so that most clays may con- 







Fig. 110. — Machine fob Blowing Window Glass. The ring, B, is filled with glass 
before the blow begins. Air is forced through the tube, W. After the cylinder 
is blown it is removed from the ring and flattened out as described in the text. 
By courtesy of The Scientific American. 



GLASS AND POTTERY 



273 



tain quartz, mica, and compounds of iron as well as calci- 
um silicate and limestone. 

By pottery is meant anything made of clay. This in- 
cludes everything from a brick to a china cup, but the 
clay used in making bricks is quite impure, while that in 
the china cup is kaolin. The poorer grades of clay are 
also used for flower pots, tile, drain pipes, and many 
other things which are classed as earthenware. On ac- 
count of the iron compounds which these contain they are 




Fig. 111. — Filter-press for Filtering Clay after Washing. During the washing 
the fine particles of clay float away from the coarse grains. The water carrying 
this floating clay is passed into a filter-press as shown in the picture and squeezed 
nearly dry. 



often red, gray, and black, according to how hard they 
have been baked, i. e., to how high a temperature they 
have been subjected. Earthenware is quite porous and 
brittle. 

China is a fine grade of porcelain, which is the best 



274 



A PRACTICAL CHEMISTRY 



kind of pottery. In making porcelain the coarser grains 
of the kaolin are removed by mixing the kaolin with 
water and floating off into a filter press all the finer por- 
tions. This very fine clay is squeezed nearly dry in the 
filter press and is then mixed with finely powdered 
quartz and feldspar or other fusible materials. Enougli 




Fig. 112. 



-Pressing Tiles from the Clay which Has been Removed from the 
Filter-press Shown in Figure 111. 



water is added to the mixture to make it dough-like. This 
dough is next shaped by putting it on a horizontal revolv- 
ing wheel and molding it with the hands as it revolves. 
Bent wires are also used to press against the revolving 
lump and thus give it shape. In some cases the clay mix- 
ture is pressed into molds. For example, in making a 
pitcher the two halves are molded and stuck together 
by pressure. The cups, plates, et cetera are then placed 
on racks to dry. When dry the ware is put into earthen- 



GLASS AND POTTERY 



275 



ware boxes, the various articles being prevented from 
touching one another. These boxes are covered and 
packed close together in a round kiln made of brickwork. 
The kiln is closed and gradually heated by furnaces 
which force in hot air. The heat is thus gradually in- 




Fig. 113. — Kilns for Burning Tiles. Similar kilns are used for burning other 
kinds of pottery. The small clay boxes seen in the picture are used to hold the 
tiles during the burning. They are called "saggers." Somewhat similar saggers 
are used for holding dishes. These are stacked in the kilns until the kilns are 
filled. The doorways are then closed and the fires started. 



creased for some days, and then the kiln is slowly cooled, 
The ware when removed from the kiln is porous, rather 
rough, and of a gray or white color. The burning partly 
fuses the material of the porcelain, leaving it translucent 
when thin. The ware must next be decorated and glazed. 
Glazing consists of covering the ware with a thin layer 
of glasslike material. The glaze is made of finely ground 
feldspar, or of soft glass. This is mixed with water to 



276 



A PRACTICAL CHEMISTRY 



about the consistency of paint. The dishes are dipped 
into this mixture, dried, and again baked or "fired," as 
it is called, thus fusing the glaze to the ware, filling up 
its pores, and making it smooth. 

The decorating is sometimes done before and some- 
times after glazing. In any case, much of it must be 




Fig. 114. — Use of Tile for Floor and Side Wall of a Hospital Kitchen. Cour- 
tesy of the Mosaic Tile Co. 



hand work — either hand painting or a sort of stenciling 
or printing. The coloring matter must all be mineral 
compounds to withstand heat. The metallic oxides pro- 
duce much the same colors as in glass. After the color 
is put on, the ware must be again baked. Sometimes 
several operations and firings are required. 

On account of shrinking many pieces are broken during 
the various bakings, which greatly increases the cost of 
the higher grades of porcelain. 






GLASS AND POTTERY 277 

Another grade of pottery intermediate between porce- 
lain and earthenware is stoneware. This includes a great 
variety of articles, such as the heavy dishes of the ordi- 
nary restaurant on the one hand, and stone butter jars 
on the other. Between these limits Ave have many dif- 
ferent kinds of cheap ware. The stoneware is thicker, 
is made of more impure materials, and is not so com- 
pletely fused as porcelain. The better grades of stone- 
ware are glazed in much the same way as porcelain, 
while such articles as the butter jars are salt glazed. This 
is accomplished by throwing common salt into the fire 
which heats the kiln when at its highest temperature. 
The salt volatilizes and the sodium reacts with the clay 
forming an artificial feldspar which fuses and runs into 
the pores. 

Enamels. — Enamels are substances closely related to 
glass and are used for coating metals. The ordinary 
agate ware is an illustration of this, likewise the coating 
on bathtubs and sinks, while a fine grade of enamel is 
often used in making jewelry. 

In making enamels the first process is the manufacture 
of a sort of glass. This glass is granulated by running 
it into water while in the molten condition. It is then 
mixed with water and ground to a fine semifluid paste. 
The metal is dipped into this, dried, and heated in a 
furnace until the glass begins to melt. After this the 
ware is carefully cooled. 

The process of making enameled ware resembles the 
glazing of porcelain. 

SUMMARY 

Glass. — Glass is a mixture of silicates, more or less transparent. 
It is made by mixing sand, sodium carbonate, lime, and 
other basic substances, and fusing the mixture in a fire clay 
pot. The properties of a given sample of glass depend 



278 A PRACTICAL CHEMISTRY 

largely upon its composition. Glass is made into shape by- 
blowing and molding. Glassware must be annealed, or 
cooled, slowly. Glass may be colored by the addition of 
metallic oxides. 

Pottery. — Pottery is anything made of clay. The purest clay is 
called kaolin. The characteristics of the different kinds of 
pottery depend upon the character of the clay used in their 
manufacture. Pottery includes earthenware, stoneware and 
porcelain. China is a fine grade of porcelain. The making 
of pottery includes molding, drying, burning, decorating, and 
glazing, all of which differ somewhat for different kinds of 
ware. 

Enamels. — Enamels are substances closely related to glass, and 
are used for coating metals. 



REVIEW QUESTIONS 

1. How does the composition of glass differ from that of 
pottery? 

2. What would be the effect of heating the glass batch very 
hot at first? 

3. What is the difference between the composition of window 
glass and flint glass? 

4. How are silicates made? 

5. To which class of pottery do the evaporating dishes used 
in the laboratory belong? 

6. Why must the pottery kiln be cooled slowly? 

7. What oxide would you use to color stoneware green? 

8. What would be the objections to an unglazed pitcher? 

9. Which fuses at a lower temperature, feldspar or clay? 
10. What is agateware? 



CHAPTER XXIII 
PAINTS 

Under this heading we shall include all those liquid 
substances which are used to cover the surface of wood 
and metals either for protection or decoration. This com- 
prises not only those substances ordinarily called paints, 
but also varnishes, lacquer, stains, and enamels. 

Materials. — All these classes of substances are made up 
of liquids carrying solid matter either suspended or dis- 
solved. The liquids most used are linseed oil, turpentine, 
and alcohols. Among the solid substances are white lead, 
numerous pigments, shellac, et cetera. 

For all ordinary purposes the best paints are made by 
grinding together white lead, linseed oil, pigment of the 
desired color, and a drier. The drier is usually a com- 
pound of lead or manganese which gives oxygen to the 
linseed oil. The oil is thus oxidized and forms a hard 
film, which holds the white lead and pigment and binds 
them fast to the painted surface. To most paints also 
some turpentine is added as a solvent, which causes them 
to spread more smoothly and evenly. The turpentine 
evaporates during the drying of the paint. 

White Lead. — White lead is basic lead carbonate. It 
is made by many different methods, most prominent 
among which is the old Dutch process. This may be de, 
scribed somewhat as follows : The floor of a room is 
covered with spent tan bark (this is pulverized bark 

279 



280 



A PRACTICAL CHEMISTRY 




Fig. 115. — White Lead Pot and Buckle. The pot is filled up to the hole, H, with 
dilute acetic acid. The lower part of the pot is made smaller so that the buckles 
may not fall into the acid. 



which has been used in tanning leather). Upon this coat- 
ing of tan bark are placed a large number of earthenware 
pots of a shape similar to that shown in the figure. These 




Fig. 116. — Making White Lead. The pots, P, stand upon tan-bark, B, which is 
spread in layers over the floors and banked against the walls. F is a flue to carry 
off the vapors and gases. 



PAINTS 



281 



pots are filled up to the holes with dilute acetic acid, 
while above the holes are a number of perforated leaden 
disks of a size that will prevent their slipping down into 
the acid in the lower part of the pots. Tanbark is also 
banked around the walls of the room. A floor of wood is 




Fig. 117. — Filling the White Lead Pots. The pots are placed upon the tan-bark 
floor and the dilute acetic acid added. They are then filled with lead buckles. 
The man at the right is putting in the buckles. 



placed over these pots, and this in turn is covered with 
tan bark and pots as described, excepting one little space 
in the middle, where a flue is made for the escape of 
gases from the tan bark below. Thus floor after floor is 
built up, till, perhaps, a stack twenty floors deep is con- 
structed. This is left undisturbed for about three months. 
During this time chemical changes occur which convert 
the metallic lead into white lead. Probably the lead and 
acetic acid with the oxygen of the air form lead acetate 



282 



A PRACTICAL CHEMISTRY 



and water; the tan bark ferments and liberates carbon 
dioxide, which reacts with the lead acetate, forming the 
basic lead carbonate of white lead. The white lead, after 
being removed from the pots, is broken up and sifted to 
take out any metallic lead which did not become cor- 



1 




\ 

<fr • ■>)'**. ;,nM9H rap - 

' ^ ; §H is 1 1 



Fig. 118. — Expressing Linseed Oil. The flaxseed is put into canvas bags whicn 
are placed one above another on iron plates as in the press at the right. The press 
then squeezes out the oil as shown by the press at the left. A very strong hydraulic 
press is required for this purpose. 

roded. It is then washed through sieves to remove large 
particles and dried. It is sold either dry or mixed with 
oil. 

By the French method white lead is made in a much 
shorter time. In this process the lead is granulated and 
put into deep tanks through which acetic acid perco- 
lates. The lead is dissolved and the solution of lead ace- 
tate thus formed is permitted to drop through tanks in 



PAINTS 283 

which numerous slats cause it to splash and present a 
large surface to the current of carbon dioxide, which is 
forced up through the tanks. Thus the same chemical 
changes are brought about as previously described, but 
the product formed is said to be of inferior quality. 

Linseed oil is an oil expressed from flax seed. The oil 
is often boiled. (See Fig. 118.) 

The pigments are quite numerous and composed of a 
great variety of substances. Prominent among these, 
however, are compounds of chromium and of lead. 

Chrome yellow is lead chromate, PbCr0 4 , that is, it is 
the lead salt of chromic acid, which has the formula 
H 2 Cr0 4 . Lead chromate is made by mixing a solution of 
any soluble lead salt with one of potassium chromate, 
K 2 Cr0 4 , or of potassium dichromate; thus: 

Pb(N0 3 ) 2 + K 2 Cr0 4 = PbO0 4 + 2KN0 3 
(Lead (Potassium (Lead (Potassium 

nitrate) chromate) chromate) nitrate) 

or, 2Pb(N0 3 ) 2 + K 2 Cr 2 7 + H 2 == 2PbCr0 4 + 2HN0 3 
(Potassium 
dichromate) +2KN0 3 

When chrome yellow is mixed with Prussian blue a 
green pigment is obtained. Another green is a basic 
oxide of chromium of the formula, Cr 2 0(OH) 4 . Paris 
green is a complex compound containing copper and ar- 
senic. It is very poisonous. 

Chrome orange and chrome red are basic lead chro- 
mates made by boiling lead chromate with an alkali, as 
sodium hydroxide. 

Among other substances used for yellow pigments, 
orpiment, the yellow sulphide of arsenic, As 2 S 3 , may be 
mentioned. Black paints often contain lampblack and 
other forms of carbon. 



284 A PRACTICAL CHEMISTRY 

White lead, as previously described, is much used in 
white paints, but they are sometimes found to contain 
lead sulphate. Zinc oxide, ZnO, is also used. Zinc oxide 
is particularly desirable for paints that are exposed to 
air containing hydrogen sulphide, since this compound 
blackens lead paints. Among red compounds for fine 
pigments vermilion perhaps is the best, but for cheaper 
and coarser paints red lead, Pb 3 4 , and Venetian red, 
Fe 2 3 , are used. Vermilion is a bright red sulphide of 
mercury of the formula HgS. It is of the same composi- 
tion as the ore of mercury, cinnabar. Vermilion is made 
by heating mercury and sulphur together and then sub- 
liming the product. 

Many cheap pigments are now made by dyeing clay 
with aniline dyes, though these are not so desirable as 
the mineral pigments. 

The all-important use of red lead is for painting the 
first coat on steel structures to protect the steel from 
rusting. Metals are sometimes protected and decorated 
by painting with metallic paints, i. e., paints containing 
powdered aluminium, brass, or bronze. 

Turpentine. — Turpentine is a thin, volatile oil made by 
distilling the gum obtained from pine trees. It is often 
adulterated with gasolene. Many thousands of gallons 
of turpentine are used each year in paints and varnishes. 

Varnishes. — Varnishes are solutions of gums and resins 
in turpentine, alcohol or oil. Shellac is a resin much used 
for this purpose, while wood alcohol is one of the most 
important solvents. Varnishes are sometimes colored but 
are often nearly transparent. Black varnishes, usually 
called Japans, contain asphaltum. 

Varnishes are used as a finishing coating on the sur- 
face of a great variety of objects. They are usually in- 
soluble in water and hence are a protection as well as 
being ornamental. Quite often for indoor work the wood 



PAINTS 



285 



is simply stained with a solution of coloring matter and 
then varnished. The grain of the wood is thus left 
visible. 

SUMMARY 

Under paints we include all liquid substances used to cover wood 

or metal for protection or decoration. 
They are liquids containing solid matter, either in suspension or 

in solution. 
Ordinary paints contain linseed oil, white lead, pigment, and 

sometimes a drier. 
The oil is obtained by expressing it from flaxseed. 
The white lead is basic lead carbonate, and is made by the old 

Dutch process, or French method, or by one of the more 

modern processes. 
Driers are substances which carry oxygen to the linseed oil. 



Pigments 



Red 

Orange 
Yellow 

Green 

Blue 
Black 

White 



Vermilion, HgS 
Red lead, Pb 3 4 
Venetian red, Fe 2 3 
Chrome red. or basic 
chromate 



lead 



Chrome orange 
r Chrome yellow, PbCr0 4 
\ Orpiment, As 2 S 3 

Chrome yellow mixed with 
Prussian blue 

Basic chromium oxide Cr 2 
(OH), 
__ Paris green 

Prussian blue 
J* Lampblack 
I Other forms of carbon 

f White lead 
■j Lead sulphate, PbSO* 
L Zinc oxide, ZnO 



286 A PRACTICAL CHEMISTRY 

Turpentine is an oil distilled from the gum obtained from pine 

trees. 
Varnishes are solutions of gums and resins in alcohol, turpentine, 

or oil. 

EEVIEW QUESTIONS 

1. What are some of the advantages gained by using paint? 

2. After drying how would a surface painted with white 
lead and linseed oil differ from one painted with white lead and 
turpentine ? 

3. What becomes of the acetic acid in the process of making 
white lead by the Dutch method? 

4. Chrome yellow is a salt of what acid? 

5. What is Prussian blue ? 

6. How could you get metallic mercury from vermilion? 

7. How can an expert tell from what kind of tree a sample 
of stained wood has been made? 

8. What is shellac? 



CHAPTER XXIV 
PHOTOGRAPHY 

Nearly everyone has perhaps noticed that plants which 
grow in the dark or those portions of plants which are 
underground are usually white. The green color of those 
portions which are in the light is clearly due to the 
chemical action of light. Various other illustrations of 
this sort may be mentioned: For example, the fading of 
wall papers and carpets and the decomposition of nitric 
acid when exposed to the light. 

The action of light upon silver compounds is the foun- 
dation of photography. This action is of several different 
sorts ; thus a pure solution of silver nitrate in a clean bot- 
tle may remain exposed to the light for any length of 
time without apparent change, but a drop of this same 
solution on the hand or on any organic matter is black- 
ened almost at once on coming into sunlight. This is due 
to the reduction of the silver nitrate by the organic mat- 
ter to free silver, which in the finely divided condition 
looks black. Silver chloride on exposure to light passes 
through a variety of shades of color due to the loss of 
chlorine. In this case the process will go on to a con- 
siderable extent without the presence of organic matter. 
A similar action occurs with silver bromide, and this is 
the silver salt used in making photographic films and 
plates. Since ruby light does not act upon silver com- 
pounds it has been found possible to do photographic 
work under such a light. In the making of films and 

287 



288 A PRACTICAL CHEMISTRY 

plates a solution of ammonium bromide and gelatine is 
mixed with silver nitrate, the bromide being in excess. 
This is kept warm for some time, and an emulsion of 
silver bromide in the gelatine is formed. The long-con- 
tinued heating renders the silver bromide more easily 
affected by light on exposure, i. e., more sensitive. The 
emulsion is cooled and thoroughly washed to remove the 
ammonium nitrate and excess of ammonium bromide. It 
is then melted and used to coat glass plates and films. 

In the camera the plate or film is exposed for a time, 
varying from a thousandth of a second to a minute, to 
the bright image produced by the lens. The picture is 
thus recorded on the plate, though invisible to the eye. 
That is, wherever the light has come in contact with the 
silver bromide it has been changed and rendered easily 
reduced to metallic silver. 

On removal from the camera the picture must undergo 
development. This is done by placing it face up in a 
bath of developing solution consisting of some reducing 
agent mixed with an alkali. The reducing agents most 
used are organic substances, such as hydroquinone and 
pyrogallic acid. The action of the developer deposits 
metallic silver in those portions of the film (or plate) 
where the light has acted, in proportion to the intensity 
of the light, i. e., where the picture has been brightest 
(for example the white portions), the most silver is de- 
posited. The result is that all highly lighted portions of 
the object are represented by heavy deposits of silver on 
the plate or film, while shadows and black parts of the 
object (which reflect little or no light) are represented by 
little or no deposit. On this account the result, being just 
the reverse of the object, is called a negative. The nega- 
tive is next put into a bath of sodium thiosulphate, 
Na 2 S 2 3 , commonly called "hypo," which dissolves out 
all silver bromide that has not been reduced by the de- 



PHOTOGRAPHY 289 

veloper, leaving those portions more or less transparent, 
while those portions where there is metallic silver are 
correspondingly opaque. The negative must be washed 
free of "hypo" with pure water and then dried. 

From the negative pictures are printed. This is usually 
done on paper, though sometimes on glass for transparen- 
cies and lantern slides. Several different kinds of pho- 
tographic paper are used. On one class of papers, a coat- 
ing of an emulsion of the same character as that on the 
plates and films is spread. This paper is placed under 
the negative in a suitable frame and exposed to light. 
The light passes through the transparent parts of the 
negative and acts upon the paper below. The paper is 
then developed and otherwise treated as a photographic 
plate. This time the metallic silver is deposited in the 
portions of the picture corresponding to the dark parts of 
the object. In another class of papers silver chloride is 
used. These papers are printed slowly in bright sunlight. 
The paper under the exposed portions gradually turns 
dark. Such papers do not require developing, but must 
be toned, i. e., the silver must be replaced with gold and 
the unchanged silver salt removed. 

The glass plates used for transparencies and slides are 
of the same character as the photographic plates, and 
are printed and developed in the same way as developing 
papers. 

Blue Prints. — Blue prints are made on a different plan. 
They contain no silver but depend upon the fact that 
ferric iron in the presence of organic matter is reduced 
by bright sunlight to ferrous, and also upon the action of 
potassium ferricyanide with ferrous and ferric salts. It 
will be recalled that the ferricyanide produces no precipi- 
tate with ferric salts and a deep blue precipitate with fer- 
rous salts. A mixture of potassium ferricyanide and fer- 
ric ammonium citrate is used to sensitize the paper. The 




Fig. 119A. — A Negative. It will be noticed that all the white parts of the negative 
are dark in the print, Figure 119B. 



§^^5^^BjL * ■' " "-*-•*•-- 










^^^K 










m, ' HJHB jKT" &' 














■-*. 






EV^9 Hte* *sl '''^ ; 










: ' "* ■' '■ ; ;-: 
























SEa** 






. •■.■■■ 




-.- 






:"*... - • 










.'■■'''. ■ ■■- ^ -i- *■' 


i- i 


:: ;" 


" , 


-■' : -'" .■'■.* '.•" 


B - ^^s*" , _ ^ 








Jlfli 



Fig. 119B. — Print, or Positive Corresponding to the Negative in Figure 119A. 



PHOTOGRAPHY 291 

ferric salt on the paper is reduced under the transparent 
portions of the negative, but is not reduced under the 
opaque portions. When the paper, after exposure, is 
put into water it turns blue where it has been acted upon 
by light, while from the other parts the chemicals wash 
away and leave the paper white. 

Blue print paper, ferriprussiate paper it is often called, 
is much used for making prints of plans for buildings and 
machinery. The various other papers need not be dis- 
cussed here. Likewise color photography has no place 
in this chapter. 

Flash Powders. — When it is desired to take pictures at 
night or in poorly lighted places, recourse is had to arti- 
ficial light produced by flash powders. These are usually 
made of powdered magnesium mixed with an oxidizing 
agent. 

SUMMARY 

The chemical action of light is illustrated by many changes 
in the objects around us. Photography is based upon the 
action of light upon silver compounds. 

In making a negative the silver salt is made into an emulsion 
with gelatine. This emulsion is used to coat a plate or film. 
After exposure to light the silver salt is reduced by the de- 
veloper. The excess of silver is dissolved in a "hypo" solu- 
tion. "Hypo" is sodium thiosulphate, Na 2 S 2 3 . Pictures 
are printed upon papers which require development, those 
which require toning, and upon blue print papers. 

Flash powders are usually composed of powdered magnesium and 
an oxidizing agent. 

REVIEW QUESTIONS 

1. Why are plants growing in a dark cellar white? 

2. How does a colorless solution of silver nitrate dye hair 
black? 



292 A PRACTICAL CHEMISTRY 

3. What test could you make for silver chloride? 

4. Why do red objects show black and violet objects white 
in a photograph? 

5. Why is the plate after developing called a negative? 

6. Why must the excess of silver be removed by hypo? 

7. What is the smoke produced by burning a flash powder? 

8. When a line drawing is printed upon blue print paper, 
what color will the lines be? 



APPENDICES 



APPENDIX A 

EXPERIMENTS FOR DEMONSTRATION AND LABORATORY 
EXERCISES 

The following experiments include both laboratory exercises 
and demonstrations. A larger number of each is offered than 
some teachers may consider necessary. This is done in order 
that the needs of all may be supplied. The teacher will use his 
own judgment as to which are best suited for laboratory work 
under the existing circumstances. 

PHYSICAL AND CHEMICAL CHANGES 

Exp. 1. — Heat a platinum wire, a bit of magnesium, and a 
stick in a Bunsen burner flame. In which of these cases is the 
composition changed? 

Exp. 2. — Drop a glass, a stone, and a torpedo. In which of 
these cases is the composition changed? Give definition of physi- 
cal change. Of chemical change. 

Exp. 3. — Heat some splints in a test tube and notice all 
changes which occur. Which of these are physical? Which 
chemical ? 

Exp. 4. — Write with silver nitrate solution upon white paper 
and place in the sunlight. What causes the chemical change? 

Exp. 5. — Moisten a filter paper with potassium iodide solution, 
place it upon a flat sheet of copper. Join the copper to the 
negative pole of a battery, and mark upon it with the positive 
wire. What is the cause of the brown marks thus formed? 

Exp. 6. — Drop a little zinc into hydrochloric acid in a test 
tube. Is the cause of the chemical change in this case the same 
as in Exps. 3, 4, or 5? 

295 



296 A PRACTICAL CHEMISTRY 

Exp. 7. — Mix some pulverized tartaric acid with some baking" 
soda. Is there a chemical change? Add water to the mixture. 
Explain the result. 

Exp. 8. — Grind together in a mortar a very little sulphur and 
potassium chlorate. What is the cause of the chemical action? 

Exp. 9. — Place a crystal of iodine on a small bit of phosphorus. 
Explain the change. Name the various agents which will bring 
about chemical changes. 



MIXTUEES AND COMPOUNDS 

Exp. 10. — Put into a heap a promiscuous collection of articles, 
as pens, matches, keys, splints, test tubes, et cetera. Notice con- 
cerning this mixture: (1) its constituents retain their properties; 
(2) they are not present in any definite quantity; (3) they can 
be separated mechanically. 

Exp. 11. — Pass an electric current through dilute sulphuric acid 
(1-20) in a Hoffman apparatus. {See Fig. 7.) Notice: (1) the 
resulting products of the decomposition have properties unlike 
those of the solution; (2) they are of definite volume; (3) they 
were not separated mechanically. Is the liquid which is being 
decomposed a compound or a mixture? What is necessary to de- 
compose a compound? 

PROPEETIES OF WATER 

Exp. 12. — Put centigrade and Fahrenheit thermometers for 
some minutes into melting ice and note the temperature on each. 
Place them also in steam over boiling water for a few minutes 
and read the temperature. Also find the temperature of the boil- 
ing water. 

Exp. 13. — Fit a test tube with a cork and long vertical de- 
livery tube. Fill the test tube and delivery tube with water at 
room temperature. Stand the test tube in ice water and observe 
the change in the volume of the water. Repeat the experiment, 
this time standing the tube in hot water. Again do the experi- 
ment, surrounding the tube with a mixture of ice and salt. 



APPENDIX A 297 

SOLUTION 

Exp. 14. — Stir with a thermometer 50 grams of ammonium 
nitrate in 20 c. c. of water in a beaker, noting- the temperature 
of the water before and after. What becomes of the ammonium 
nitrate? Why the change in the temperature? (N. B. — Keep 
this and the following- solutions for future use.) 

Exp. 15. — Heat a flask of water to boiling- and add mercuric 
chloride so long as any will dissolve. Try to dissolve the same 
amount of the chloride in an equal volume of cold water. What 
part does the heat play? (Set the hot solution aside where it 
will not be disturbed.) 

Exp. 16. — Try to dissolve a gram of camphor in 10 c. c. of 
alcohol. Result? In 10 c. c. of water. Result? 

Exp. 17. — Let a little vichy escape from a siphon into a beaker. 
Why does the water effervesce? 

PUKIFICATION OF WATER 

Exp. 18. — Completely evaporate some drinking water in a 
clean porcelain dish on a water bath and notice whether any 
sediment remains. 

Exp. 19. — Examine under a high-power microscope a sample 
of water which has been contaminated by allowing decayed meat 
to lie in it for some days. Boil some of this contaminated water 
for several minutes and again examine under the microscope. 
Result? What is gained by boiling drinking water? 

Exp. 20. — Gently warm some water in a flask and notice the 
gas bubbles which soon begin to form on the walls of the flask. 
Explain why these bubbles appear. Why does boiled water have 
a different taste from that which has not been boiled? 

Exp. 21. — Color some water by the addition of a few drops 
of ink. Distil the water from a retort and examine the distillate 
for impurities. Repeat, using water containing kerosene oil. 
Is the oil removed from the water? What kinds of substances 
can be removed by distillation? 

Study 1. — Make a careful study of one or more dia- 
grams of water filters for city water supply and copy one 
diagram, 



298 A PRACTICAL CHEMISTRY 



CRYSTALLIZATION 

Exp. 22w From Solution. — Let the solution of ammonium ni- 
trate made in Exp. 14 stand exposed to the air in an open dish 
for some days and examine the crystals formed. Observe the 
formation of crystals in the solution of mercuric chloride of 
Exp. 15 as it cools. Melt some crystals of sodium thiosulphate 
in a large test tube and set the tube aside to cool, thus forming 
a supersaturated solution. When this solution is cold add a small 
crystal of sodium thiosulphate and observe results. Explain 
why the temperature of the solution changes. 

Exp. 23. From a Fused Mass. — Melt a quantity of sulphur 
in a dish and then let it cool. When a crust forms over the sur- 
face of the sulphur make a hole through this crust and pour out 
the molten sulphur. Examine the crystals which remain under 
the crust. 

Exp. 24. From a Vapor. — In a large, clean, dry flask heat 
some iodine and notice the crystallization of the vapor on the 
cool portion of the flask. 

Exp. 25. Water of Crystallization. — Heat a crystal of copper 
sulphate in a test tube, holding the tube with its mouth inclined 
downward. Notice all changes which occur. Where does the 
water come from? Is the change in the crystal chemical or 
physical ? 

Exp. 26. — Examine some bright, clean crystals of sodium sul- 
phite and expose them to the air on a glass plate for an hour. 
Explain the change in the crystals. What word describes this 
process? 

Exp. 27. — Repeat Exp. 26, using some fused sodium hydroxide. 
By what name is the change in the sodium hydroxide called? 
Explain the change. 

COMPOSITION OF WATER 

Exp. 28. — Repeat Exp. 11, but this time carefully measure 
the volume of each gas; collect some of each in a test tube, and 
test whether it will burn and support combustion. 



APPENDIX A 299 



OXYGEN 

Exp. 29. — Fold a little mercuric oxide (the red powder formed 
by heating mercury in the air) in a strip of paper and pour it 
to the bottom of a small test tube without permitting any of 
the oxide to adhere to the walls of the tube. Heat the tube care- 
fully and from time to time introduce a spark on the end of a 
splint. Note results. What is the coating on the walls of the 
tube? 

Exp. 30. — Heat a little potassium chlorate in a test tube, no- 
ticing all changes which occur. When the melted chlorate bubbles 
violently hold the tube well away from the face and drop into 
it a lighted splint. Explain. 

Exp. 31. — Generate and collect several jars of oxygen by heat- 
ing a mixture of one part of manganese dioxide to three parts 
of potassium chlorate, using apparatus as shown in Figure 8. 
Using a deflagrating spoon, burn in one jar some charcoal, in a 
second some sulphur and in the third some red phosphorus. Com- 
pare the combustion in each case with the combustion of the same 
substance in the air. Observe also the character of the products 
of the combustion. 

Exp. 32. — Pour about 200 c. c. of water into a large acid bottle 
and add to the water about 15 grams of sodium peroxide. Make 
a spiral of a piece of fine iron wire and wrap one end of this 
wire around the tip of a match. Light the match tip in a flame 
and thrust the wire into the acid bottle. What is the product 
of the combustion of the iron wire? 

Exp. 33. — Using the potassium chlorate mixture, a copper re- 
tort, and gas holder, generate and collect a considerable quantity 
of oxygen. Arrange an ozonizer composed of a glass tube carry- 
ing a coil of wire within and surrounded by another coil of wire, 
as shown in Figure 10. Join the outer coil to one terminal of 
an induction coil and the inner coil to the other terminal. Turn 
on the electric current and pass a stream of dry oxygen through 
the ozonizer tube. The oxygen may be dried by bubbling it 
through concentrated sulphuric acid. Notice the peculiar odor 
of the ozone as it issues from the ozonizer tube. Notice its 




300 A PRACTICAL CHEMISTRY 

oxidizing power by bringing it in contact with moist lead sul- 
phide, black rubber and other oxidizable substances. 

HYDROGEN 

Exp. 34. — Arrange a Kipp apparatus (see Fig. 15), putting 
into the middle bulb two or three hundred grams of zinc and fill- 
ing the upper and half of the lower bulbs with dilute hydrochloric 
acid. Pass the delivery tube of the generator up into a "hydro- 
gen gun," closing the upper hole of the gun with the thumb. 
When the gun is probably filled with hydro- 
gen close the stopcock of the Kipp, and, re- 
moving the thumb, bring a flame to the upper 
hole of the gun. How do you explain the 
explosion of the hydrogen? If the explosion 
is delayed, how do you explain the delay? 
What is the source of the hydrogen used in 
this experiment? 
Fig. 120.— The Hydho- Exp. 35. — Avoid an explosion when light- 

fs E he?d U with T mouth in S a hydrogen jet thus: Pass the delivery 
down while being tube up into an inverted test tube and allow 
filled, the touch-hole, t ^ e hydrogen to flow into the test tube until 
meanwhiie eP with°the ii} is probably filled. Holding this test tube 
thumb. vertical and with its mouth down, bring it 

to a flame which is kept at a safe distance, 
thus lighting the hydrogen in the tube. If the hydrogen burns 
quietly in the test tube, it may be carried back to the jet while 
still burning and used to light the jet. If this method is em- 
ployed, many dangerous accidents may be avoided. 

Exp. 36. Water Is the Product of the Combustion of Hydrogen. 
— Thoroughly dry a stream of hydrogen by passing it over fused 
calcium chloride and light it as described above. Over this burn- 
ing jet of hydrogen hold a funnel, the stem of which passes 
into a large cold glass bulb. (See Fig. 15.) Water will collect 
on the walls of this bulb. Why? 

Exp. 37. Hydrogen is lighter than air.— Fill a jar with hy- 
drogen and, keeping it mouth downward, try to pour the hydro- 
gen upward into another jar, as shown in the illustration. (See 



APPENDIX A 301 

Fig. 12.) Afterward bring the mouth of each jar to the Bunsen 
flames. 

Exp. 38. Combustion in Hydrogen. — Light a long splint and 
pass it upward into a jar of hydrogen, withdraw the splint and 
repeat several times until you have observed all that occurs. Ex- 
plain results. 

Exp. 39. — Hydrogen as a reducing agent may be illustrated by 
putting a few pieces of copper oxide into a clean, dry delivery 
tube about 8 inches long, attaching this to the delivery tube 
of a hydrogen generator, and, while the hydrogen (free from 
air) is passing freely, heating the oxide in a Bunsen flame. 
Examine the substances left in the tube and explain the ac- 
tion. 

Exp. 40. Decomposition of Water. — Fill a piece of half -inch 
gas pipe about three feet long with bright iron tacks, fit one end 
of this pipe with a delivery tube and attach the other end of the 
pipe to a flask half filled with water. Heat the gas pipe in a 
combustion furnace and at the same time bring the water to a 
boil. {See Fig. 16.) Examine the gas which escapes from the 
delivery tube for hydrogen. What has become of the oxygen 
of the water? 

Exp. 41. Action of Sodium on Water. — Into a test tube con- 
taining about a half inch of water at the bottom drop one at 
a time a number of small pieces of sodium. Bring a lighted 
match to the mouth of the tube. Why does the hydrogen thus 
obtained burn with a yellow flame? Boil off the water and ex- 
amine any residue which may remain. 

Exp. 42. To Make Hydrogen Peroxide. — Add about 10 c. c. 
of concentrated phosphoric acid to ten times its volume of water 
and cool the mixture. Slowly stir into this a teaspoonful of 
barium dioxide and let the precipitate settle. Hydrogen peroxide 
is found in solution in the liquid. 

Exp. 43. Hydrogen Peroxide as an Oxidizing Agent. — Add 
some hydrogen peroxide to some precipitated lead sulphide; the 
change in color of the sulphide shows that it has been oxidized. 
Pour some hydrogen peroxide as obtained at the drugstore on 
blood, pus, or other easily oxidized substances, and observe re- 
sults. 



302 A PRACTICAL CHEMISTRY 



ACIDS AND BASES 

Exp. 44. — Make phosphorus pentoxide by burning a bit of 
phosphorus in a jar of air containing a few cubic centimeters of 
water. Close the mouth of the jar and shake the water and oxide 
together till the latter is dissolved. This solution is phosphoric 
acid. Try its action on red and blue litmus paper and on baking 
soda. Taste it. Record results. 

Exp. 45. — Burn a bit of sodium in an evaporating dish and 
dissolve the sodium oxide in a few drops of water. The resulting 
solution of sodium hydroxide is a base; try its action on red and 
blue litmus paper. Rub a drop of the solution between the 
fingers. Dilute and taste a little of the solution. Record results. 

Exp. 46. — In the same way make and dissolve other oxides, 
as oxides of carbon, sulphur, magnesium, and obtain calcium and 
barium oxides from the instructor. Treat these oxides as in 
Exps. 44 and 45. Tabulate results and indicate which solutions 
are acids and which bases. 

Exp. 47. Neutralization and Salt Formation. — Dissolve about 
a gram of sodium hydroxide in an evaporating dish and add 
dilute hydrochloric acid a little at a time with continuous stirring, 
till a drop of the mixture when removed with a stirring rod and 
placed on red and on blue litmus will have no effect on either. 
Taste the result. Evaporate nearly to dryness and let cool; if 
crystals do not form evaporate further. (The solution may be 
allowed to stand and evaporate spontaneously.) Examine the 
crystals. What are they? To what class of substances do they 
belong? How do the properties of this substance compare with 
those of sodium hydroxide and of hydrochloric acid? Why have 
these substances lost their properties? 

Exp. 48. Neutralization Is Quantitative. — Clamp two clean 
burettes perpendicularly as shown in Figure 18. Fill one with 
dilute hydrochloric acid and the other with a solution of sodium 
hydroxide. Measure from the acid burette 10 c. c. of acid into a 
small beaker, add two drops of phenolphthalein solution, and 
then carefully run in the sodium hydroxide with continuous stir- 
ring till a faint permanent pink tint is obtained. Carefully note 



APPENDIX A 303 

the volume of sodium hydroxide used. Repeat several times and 
compare the volumes of hydroxide used. 

Exp. 49. Electrolysis of Brine. — Fill a U tube nearly full of 
ten per cent, salt solution to which has been added a little neutral 
litmus solution. (See Fig. 19.) Place a platinum electrode in 
each tube and connect with a source of electric current. Observe 
all changes in the contents of the tube and examine carefully the 
gases which escape from them. Notice also which tube contains 
the positive and which the negative electrode. Read your text 
in connection with the experiment; what large commercial enter- 
prise is illustrated by this experiment? Make a labeled diagram 
of apparatus for this industry. Name all products formed in 
the electrolysis of brine. 

Exp. 50. Preparation of Chlorine (Work under Hood). — Ar- 
range apparatus consisting of flask with a rubber stopper carry- 
ing a thistle tube and long glass delivery tube as shown in Fig- 
ure 22. Support the flask by a clamp on an iron stand and pro- 
tect it from the flame by a wire gauze. Fill the flask one-third 
full of manganese dioxide and add enough concentrated hydro- 
chloric acid to make the mixture into a thin paste when shaken. 
Gently warm the flask and collect the escaping chlorine by dis- 
placement of air in glass bottles or cylinders having their mouths 
nearly closed by glass plates over the lower side of which has 
been smeared a little vaseline. When each jar is filled with the 
gas it may be completely sealed by covering it with a vaseline- 
coated glass plate. From what compound is the chlorine obtained 
in this experiment 1 ? Put a little potassium permanganate, which 
is a compound rich in oxygen, into a test tube and add a few 
drops of hydrochloric acid. Do you obtain chlorine from this 
mixture? Repeat, using a little potassium chlorate. What is 
the action of the manganese dioxide, potassium permanganate, 
and potassium chlorate in these reactions? 

Exp. 51. Properties of Chlorine. — Notice the color and odor 
of chlorine. Put into a jar of chlorine strips of wet and dry 
calico, printed paper, and writing, and observe the effect on each. 
Into a jar of chlorine lower a jet of burning hydrogen and 
notice result. Thrust a burning wax taper into a jar of chlorine. 
Wax is composed of carbon and hydrogen. Explain the com- 



304 A PRACTICAL CHEMISTRY 

bustion of the taper in chlorine. Heat an iron wire red hot and 
thrust it into a jar of chlorine. What is the product? Repeat, 
using a copper wire. Lower some burning sodium into chlorine. 
What is the product in this case? Sprinkle some finely pow- 
dered metallic antimony into a tall jar of chlorine. What would 
you conclude concerning the action of chlorine on metals'? 

Pass chlorine into cold water; is it soluble? Pass chlorine 
into slaked lime; what industry is based on this reaction? (See 
text.) 

Exp. 52. Action of Bleaching Powder. — Make a mixture of 
bleaching powder and water, and with this mixture saturate 
strips of calico or other fabrics which you have found may be 
bleached by chlorine; then dip them into dilute hydrochloric 
acid. Result? Make labeled diagram of the industry repre- 
sented by this experiment. 

Exp. 53. Preparation and Properties of Hydrochloric Acid. 
■ — Arrange apparatus as for chlorine. Put into the flask about 
100 grams of common salt and enough water to equal the salt 
in volume. Add concentrated sulphuric acid in small quantities 
through the thistle tube. Collect a jar of the gas as you did 
chlorine. Blow your breath across the mouth of this jar. Ex- 
plain result. Bring a bottle of ammonium hydroxide to the 
mouth of a bottle of hydrochloric acid gas. Read text and ex- 
plain results. By the aid of a short rubber tube hang a funnel 
onto the end of the delivery tube and place the funnel in a dish 
of water. Pass the hydrochloric acid gas thus into the water 
for some time. Try the action of this solution on litmus paper. 
How does the hydrochloric acid of the laboratory differ from 
your solution? Boil some of your solution in an evaporating 
dish till it is much reduced in volume. Does the solution change 
in concentration? 

Exp. 54. Test for Hydrochloric Acid and Chlorides. — To a solu- 
tion containing hydrochloric acid or soluble chlorides add a few 
drops of a solution of silver nitrate; a precipitate of silver chlo- 
ride is obtained. Divide this precipitate into two parts ; one part 
is found to be insoluble in nitric acid, and the other portion 
dissolves freely in ammonium hydroxide. Test drinking water 
in this way for chlorides. 



APPENDIX A 305 

Exp. 55. Tli© Barometer. — Using the blast lamp, carefully 
fuse and close one end of a glass tube about 35 inches long hav- 
ing an internal diameter of about 3-16 of an inch. Fill this tube 
with clean mercury, close the open end with the thumb and invert 
the tube, placing its open end in a cup of mercury, and care- 
fully remove your thumb without permitting any air to enter. 
{See Fig. 24.) Clamp the tube in a perpendicular position. Why 
does the mercury stand up in the tube? Measure the height of 
the mercury in inches and in centimeters. Compare this height 
with the barometer reading for the same time and place. 

Exp. 56. Wet and Dry Bulb Thermometers. — Choose two ther- 
mometers which are of equal reading and wrap the bulb of one 
with a lamp wick. Clamp the two thermometers in a burette 
stand and permit the lampwick to extend down into a bottle of 
water. After the thermometers have stood for an hour or so 
take the reading of each. Account for the difference in the read- 
ings. Read the thermometers on several successive days and ex- 
plain why the difference between them should not always be 
the same. 

Exp. 57. Volume of Oxygen in the Air. — Invert a measuring 
cylinder of air in a deep, wide cylinder of water, and clamp it 
so that the water within and without will be at the same height. 
Read the volume of air in the cylinder. Twist a wire about a 
small stick of phosphorus, and thrust the phosphorus up into the 
cylinder till it stands several inches above the water. {See Fig. 
25.) Let the apparatus stand for twenty-four hours and again 
read the volume of air in the cylinder. How do you account for 
the loss in volume? Does this give you an exact measure of any 
constituent of the air? 

Exp. 58. Nitrogen from Amm onium Nitrite. — Arrange ap- 
paratus as for hydrogen. Put into the flask about 15 grams of 
sodium nitrite and an equal weight of ammonium chloride; add 
enough water to dissolve these salts. Gently warm the flask and 
collect the gas in jars over water. Thrust a burning splinter into 
one jar and burning sulphur into another. Results? Has the 
gas an odor? Does it have any effect upon litmus? 

Exp. 59. A Test for Carbon Dioxide. — Make some carbon di- 
oxide by burning a bit of charcoal in a deflagrating spoon in a 



306 A PRACTICAL CHEMISTRY 

jar of air. Pour in a few c. e. of lime water; close the mouth 
of the jar and shake. A white, milk-like precipitate of calcium 
carbonate shows that carbon dioxide was in the jar. If the lime 
water is allowed to stand in contact with the carbon dioxide with- 
out shaking, a crust will form on the surface of the lime water. 
Try this by putting the lime water into a jar and burning the 
charcoal above it. 

Exp. 60. Carbon Dioxide in the Air. — Let a watch glass of 
lime water stand in the air for a half hour, and observe the crust 
forming upon it. 

Exp. 61. Carbon Dioxide a Product of Combustion. — Put into 
a flask a little lime water and thrust in a burning splinter, hold- 
ing it in the flask till extinguished. (See Fig. 26.) Shake the 
flask. Result? 

Exp. 62. Carbon Dioxide a Product of Respiration. — With the 
aid of a thistle tube blow your breath into lime water. Explain 
result. In how many ways can you account for the carbon 
dioxide of the air? 

Exp. 63. Water a Product of Combustion. — Hold a large, 
clean, dry funnel inverted for a moment over a Bunsen flame, 
and observe any change in the appearance of the funnel. 

Exp. 64. Forms of Carbon. — Hold a glass tube in the upper 
part of an ordinary gas flame or of a Bunsen flame when the 
holes are closed. What form of carbon is deposited on the glass 
tube? Why did it deposit? Put into a test tube some small bits 
of dry wood and carefully heat till all vapor has passed off. 
What form of carbon is left in the tube? How do these forms 
of carbon differ? Compare them with all obtainable specimens 
of various forms of carbon. 

Exp. 65. Charcoal Filter. — Make a solution of molasses, or 
brown sugar, or of a little ink in water. Put about two teaspoon- 
fuls of animal charcoal into a small beaker half filled with the 
colored solution. Bring the mixture to a boil and filter through 
a filter paper. If the color is not entirely removed, repeat. 

Exp. 66. Charcoal as a Reducing Agent. — Mix a half gram of 
powdered cupric oxide with twice its volume of powdered wood 
charcoal. Put the mixture into a small bulb tube and slip a 
rubber delivery tube, over the end of it. (See Fig. 121.) Pass 



APPENDIX A 



307 




Fig. 121. — Reduction bt 
Charcoal,. The copper ox- 
ide and charcoal are put into 
the bulb tube, B, to which is 
joined the rubber tubing, R. 
This dips into the lime water 
in L. 



the delivery tube into a test tube half filled with lime water, and 
heat the bulb for a few minutes. What gas is given off? Exam- 
ine the contents of the bulb tube. How do you explain the re- 
action. 

Exp. 67. Distillation of Coal. — Put about two hundred grams 
of soft coal into an iron tube ten inches long, closed at one end 
and the other fitted with a delivery 
tube which passes through two wash 
bottles containing water, as shown in 
Figure 128. Heat the iron tube to a 
high temperature with a burner or 
blast lamp. Examine the gas which 
escapes from the delivery tube. Will 
it burn? What do you find in the 
wash bottles? Examine the product 
remaining in the iron tube. 

Exp. 68. Marsh Gas. — Clean the ap- 
paratus used in Exp. 67, and put into 
the iron tube about a hundred grams 
of a mixture composed of equal quantities of fused sodium ace- 
tate, sodium hydroxide, and slaked lime. Heat the iron tube as 
before and collect the gas given off in glass jars. Is this gas 
combustible? Will it support combustion. Does it change lit- 
mus? Do you find anything in the wash bottles? 

Exp. 69. Acetylene. — Drop a bit of calcium carbide into a test 
tube; add a few drops of water. Light the acetylene as it comes 
from the open mouth of the test tube. Why is the flame smoky? 
Blow the flame carefully ; what change is produced and why ? 

Exp. 70. Acetylene as an Illuminant. — Using an acetylene 
generator and burner as shown in Figure 30, generate and burn 
acetylene. Examine the burner, and explain why the acetylene 
flame from the burner is unlike that from the test tube in Exp. 69. 

Exp. 71. Structure of Flames. — Light a candle and hold a 
paper down on the flame horizontally until it begins to scorch. 
Repeat, holding the paper in a perpendicular position against the 
side of the flame. What do these markings teach as to the struc- 
ture of the flame? Make a complete diagram of a candle flame. 
Light the Bunsen flame and quickly plunge a match into the inner 



308 A PRACTICAL CHEMISTRY 

blue cone. What does this show? Hold the Bunsen on one side 
and pour powdered charcoal into the flame. ~ Explain result. In 
what other ways can the same result be obtained? 

Exp. 72. Wire Gauze and Flames. — Hold a wire gauze an 
inch above a Bunsen burner, turn on the gas and light it above 
the gauze. (See Fig. 39.) Explain why it does not light below 
the gauze. Light the burner and hold the gauze down on the 
flame. Why does the flame strike through so soon? Would a 
smaller mesh give a different result? What practical application 
has been made of the principle illustrated in this experiment ? 

Exp. 73. Oxidizing and Reducing Flames. — Having closed the 
holes at the bottom of the Bunsen burner, and turned the flame 
down to about one and one-quarter inches, place the blowpipe tip 
within this flame and blow so that a blue cone is formed. This 
is an oxidizing flame. (See Fig. 35.) Direct an oxidizing flame 
against a little zinc placed in a small hollow near the end of a 
stick of charcoal, and observe the effect. Make a reducing flame 
in the same way, except that the blowpipe tip is held just out- 
side of the flame and the air blast is blown more gently, so that 
the cone formed is yellow. Use the reducing flame as you did 
the oxidizing flame, directing it against some lead oxide. (Note. 
— In all blowpipe work care should be taken to keep the charcoal 
in such a position that the flame will blow the long way of the 
coal; also to keep the fingers out of the way of the flame.) Ex- 
plain the action of the oxidizing and reducing flames. 

Exp. 74. Chemical Fire Extinguishers. — Carefully examine a 
fire extinguisher. (See Fig. 41.) Invert it and observe its action, 
directing its stream on a small fire of paper. Read directions for 
recharging, and refill the extinguisher. Arrange apparatus as 
for chlorine, put into the flask a few hundred grams of marble, 
cover the marble with water, and add hydrochloric acid as needed. 
Collect a large jar of carbon dioxide, place a lighted candle in 
a large beaker, and pour into the beaker the jar of carbon 
dioxide Result? Explain the action of the fire extinguisher, 
telling why it must be inverted, and how the carbon dioxide is 
formed. 

Exp. 75. Phosphorus. — Carefully examine specimens of red and 
waxy phosphorus. Dissolve a little waxy phosphorus in carbon 



APPENDIX A 309 

bisulphide and pour the solution on a sheet of paper. Place the 
paper on a board to dry. Explain. Can you do the same with 
red phosphorus? Drop a crystal of iodine upon a very small bit 
of phosphorus and observe results. What is spontaneous com- 
bustion ? 

Exp. 76. Phosphene. — Arrange a 50 c. c. flask fitted with a one 
hole rubber stopper carrying a long delivery tube. Let the lower 
end of the delivery tube rest in a pan of water. Place in the 
flask a bit of phosphorus as large as a pea, a stick of sodium 
hydroxide an inch long and 25 c. c. of water. Carefully bring 
the liquid in the flask to a boil and observe the bubbles of phos- 
phene as they light over the surface of the pan. If the air is still, 
good smoke rings will form. What is the composition of the 
smoke formed? 

Exp. 77. Preparation of Phosphoric Acid and Test for Phos- 
phates and Phosphoric Acid. — Burn a little phosphorus in a jar 
of air containing about 25 c. c. of water. Close the mouth of 
the jar and shake, dissolving the phosphorus pentoxide in the 
water and thus forming phosphoric acid. To a little of this solu- 
tion in a test tube add about one-fourth its volume of ammonium 
molybdate solution and gently warm the tube. A yellow precipi- 
tate of ammonium phospho-molybdate shows the presence of 
phosphoric acid or a phosphate. 

Exp. 78. Phosphate in Bone. — Burn a bit of bone in the oxi- 
dizing flame to a white ash. Dissolve some of the white ash 
(avoiding any black, unburnt portions) in nitric acid, and test 
for phosphates as given above. 

Exp. 79. Forms and Properties of Sulphur. — Slowly melt a quan- 
tity of roll sulphur in a test tube, observing carefully all changes 
which it passes through. When the sulphur has nearly reached 
the boiling point pour it quickly into a pan of cold water. Exam- 
ine the plastic sulphur thus formed and then put it into your desk 
for some days, noticing what changes occur in it from time to 
time. Dissolve some roll sulphur in carbon bisulphide (keep CS 2 
away from flames), and let the solution evaporate from a watch 
glass. Examine the crystals thus formed with a magnifying glass. 

Melt a quantity of sulphur and pour it into a folded filter. 
When a crust begins to form on the surface of the sulphur break 



310 A PRACTICAL CHEMISTRY 

through this crust and pour out the excess of sulphur. Examine 
the remaining crystals. 

Exp. 80. Sulphides. — Make a mixture of flowers of sulphur 
and iron by hydrogen (or iron filings) in the proportion of six 
grams of sulphur to ten of iron. Charge the mixture into a dry 
test tube and heat the bottom of the tube till it begins to glow. 
Watch the glow as it spreads through the mass in the tube. Ex- 
plain the reaction. What compound is formed? 

Heat a little sulphur at the bottom of a test tube till it vapor- 
izes, and drop into the tube a strip of bright copper foil. What 
is the product? Make a mixture of equal volumes of flowers of 
sulphur and powdered zinc. Pour into a conical pile on an iron 
plate or stone and with a long taper light the top of the cone. 
Examine the zinc sulphide thus formed. 

Exp. 81. Hydrosulphuric Acid (Work under Hood). — Into a 
test tube or flask fitted with a delivery tube put some ferrous 
sulphide and enough sulphuric acid (diluted, one part acid to ten 
of water) to cover the sulphide. Notice the odor. Will the gas 
burn? Pass the gas into solutions of salts of copper, antimony 
and arsenic, acidulated with hydrochloric acid. What are the 
products ? 

Exp. 82. Sulphur Dioxide. — Burn a little sulphur in the air 
and notice the odor. 

Put a few bits of copper and a little sulphuric acid into a test 
tube and heat till a reaction begins. Does the evolved gas have 
the same odor as that obtained by burning sulphur? Arrange 
apparatus as for hydrogen, put into the flask about 50 grams of 
sodium sulphite and enough water to cover it, then add sulphuric 
acid a little at a time as needed. Notice the odor of the evolved 
gas. Will this gas burn? Pass the gas into a solution of potas- 
sium permanganate till the color is destroyed. Save this solution. 
Will this gas destroy the color of other substances? (Try 
flowers, calico, straw, et cetera, both dry and damp.) Pass the 
gas into water till the solution is saturated. Try the solution 
with litmus. What does it contain? 

Exp. 83. Test for Sulphuric Acid and Sulphates, — To a little 
dilute sulphuric acid add a few drops of a solution of barium 
chloride. A white precipitate of barium sulphate, insoluble in 



APPENDIX A 311 

acids (except concentrated sulphuric) is formed and is a test for 
sulphuric acid and sulphate. Test the solution saved from Exp. 
82, also a sample of drinking water, for sulphates. 

Exp. 84. Flame Tests. — Clean a platinum wire by dipping it 
into a little clean hydrochloric acid and then burning it in the 
Bunsen flame, repeating the operation till the wire no longer gives 
color to the flame. Touch the wire to some sodium chloride and 
hold it again in the flame; observe the color of the flame. View 
the flame through a thick piece of cobalt glass. How do you 
explain the change in appearance? Note. — Any other sodium 
salt will give a similar flame. Can this be used as a test for so- 
dium compounds'? 

Clean the wire and touch it to a potassium compound, return 
it to the flame and notice the color. View this flame through the 
cobalt glass. Result? Put both sodium and potassium com- 
pounds on the wire. What color is the flame now? What color 
is it when viewed through the cobalt glass? How can you detect 
potassium in the presence of sodium? 

Exp. 85. Fertilizers — Ashes. — Examine a sample of wood 
ashes for potassium compounds as described above, also for phos- 
phates. 

Exp. 86. Nitric Acid from a Nitrate. — Arrange an apparatus 
consisting of a tubulated retort supported by an iron stand and 
extending into a flask resting in a pan of cold water. (See Fig. 
48.) Put into the retort about 25 grams of sodium nitrate (Chili 
saltpeter), and with the aid of a funnel add enough concentrated 
sulphuric acid to cover the nitrate. Slowly heat the retort till 
the acid begins to distil over, being careful that the temperature 
does not become high enough to decompose the acid and convert 
it into colored gases. When a teaspoonful of acid has collected 
discontinue the distillation and pour the contents of the retort 
into a sink into which the water is running. Wash the retort 
as soon as cool enough. Save the nitric acid for a future ex- 
periment. 

Exp. 87. Nitric Acid from tlie Air. — Arrange apparatus as 
shown in Figure 56 so that a current of air slowly passes between 
two electrodes between which an electric arc is playing, and then 
bubbles through a little water. Continue this process for twenty 



312 



A PRACTICAL CHEMISTRY 



minutes and then test the water for nitric acid, as described in 
Exp. 88. Read text and explain reactions for making the acid 
thus. 

Exp. 88. Test for Nitric Acid. — Fill a test tube about one- 
third full of the solution to be tested, drop into this a good sized 
crystal of ferrous sulphate and, holding the tube slanting, pour 
in enough concentrated sulphuric acid (letting it run down the 
side of the tube) to cover the crystal. If a nitrate or nitric acid 
is present, a brown ring will form at the surface of the sulphuric 
acid. Test thus a sample of sodium nitrate, also some of the acid 
of Exp. 87. 

Exp. 89. Some Reactions of Nitric Acid. — Put a bit of cop- 
per into one test tube and a little tin into another; add a few 

drops of nitric acid to each. Ex- 
amine all products. In what re- 
spects do the two reactions seem to 
differ? In evaporating dishes put 
strips of cotton and wool with a 
little nitric acid. Is the reaction 
the same with both? Put the nitric 
acid made in Exp. 86 into a test 
tube and place a little ball of loosely 
wound wool in the mouth of the 
tube. Heat the acid to boiling. 
Explain results. 

Exp. 90. Ammonia. — Make a 
mixture of one part ammonium 
chloride and three parts dry slaked 
lime. Fill a test tube two-thirds full with this mixture, and close 
it with a rubber stopper carrying a straight delivery tube. Collect 
a test tube full of the escaping gas by upward displacement (as in 
Fig. 122) ; close the mouth of the tube with your thumb and invert 
it in a dish of water ; then remove your thumb. Why does the water 
rise in the tube? What is formed in the tube? Hold the mouth of 
the hydrochloric acid bottle to the delivery tube. Explain the re- 
sult. What effect does the gas have on litmus paper? Put a little 
of an ammonium salt into a test tube and add a little of a solution 
of sodium hydroxide; then warm the tube. Smell the escaping 




Fig. 122. — Collecting a Test 
Tube Full of Ammonia. The 
mixture of ammonium chloride 
and lime is heated in the tube, 
A, while the ammonia collects 
in T. On removing T, it must 
be closed with the thumb. 



APPENDIX A 313 

gas; what is it? What test can you offer for ammonium com- 
pounds? Further test the solubility of ammonia in water by 
making the "ammonia fountain." Invert a large strong flask and 
clamp it in an iron stand or other suitable support. Fit the 
flask with a two-holed rubber stopper, through one hole of which 
a glass tube extends up into the flask about two inches. Connect 
a rubber tube with the outer end of this glass tube and make it 
of such a length that it will reach to the bottom of a large jar 
of water to which a few drops of phenolphthalein have been 
added. Place this jar below the inverted flask. Remove the rub- 
ber stopper. Thrust a medicine dropper filled with water through 
the other hole of the rubber stopper. Lay this part of the ap- 
paratus aside ready for use. Fit a small flask partly filled with 
strong ammonium hydroxide with a cork and delivery tube that 
will reach nearly to the top of the flask. Boil the hydroxide till 
the large flask has been filled with ammonia. (How can you tell 
when this is accomplished?) Quickly remove the delivery tube 
and replace the rubber stopper. Squirt the water from the medi- 
cine dropper up into the ammonia in the flask. The water from 
the jar will rush up through the delivery tube into the ammonia 
flask. Explain why. Why does the phenolphthalien turn 
red? Does the ammonia flask become entirely filled with water? 
Why? 

Exp. 91. Preparation and Properties of Nitrous Oxide. 
— Arrange apparatus as for oxygen, put into the test tube enough 
ammonium nitrate to fill it one-third full. Very carefully heat 
the tube and collect the resulting gas over hot water. (Why 
hot?) Has the gas any odor? Will it burn? Can you burn 
wood, iron and sulphur in it as you did in oxygen? 

Exp. 92. Nitric Oxide and Nitrogen Peroxide.— Use apparatus 
as for hydrogen. Put into the flask about ten grams of scraps 
of copper and cover them with water. Through the thistle tube 
add nitric acid a little at a time as needed. Collect the nitric 
oxide over water. Is the gas in the jar the same color as that 
in the flask? If not, why not? Collect a jar nearly full of nitric 
oxide, close the mouth of the jar with your hand, turn the jar 
mouth up and remove your hand for a moment. What is formed 
in the jar? Close the mouth of the jar again with your hand and 



314 



A PRACTICAL CHEMISTRY 



shake the jar. What evidence do you have that the gas dis- 
solves? Test the liquid with litmus. 

Exp. 93. Nitrogen Iodide. — Put a few drops of concentrated 

ammonium hydroxide upon a few crystals of iodine in a small 

crucible and gently warm for a few minutes, then pour onto 

filter paper and let stand till dry. When dry handle with care, 

observing how slight a blow is required to explode the product. 

Exp. 94. Preparation of Starch. — Wash and pare a potato and 

cut it into small pieces; grind these to a pulp with water in a 

mortar. Add a quantity of water and float the starch out from 

the other parts of the potato. Dry part of this starch and exam- 

— ^ ine under the - microscope. Repeat 

|S the experiment, using cornmeal in 

S^ 1 *^^ place of potato. Compare the 

two kinds of starch under the 
microscope. 

Exp. 95. Starch Paste. — Grind 
a little starch in a spoonful of 
cold water in a mortar, and pour 
this mixture into a half beaker of 
boiling water. Examine the starch 
paste and reserve for next experi- 
ment. 

Exp. 96. Test for Starch. — Add 
a few drops of a solution of iodine 
to some starch paste and observe 
the color. 

Exp. 97. Detection of Glucose. 

— Dissolve about one gram of 

glucose in a half test tube of 

water, add half this volume of Fehling's solution and heat to 

boiling. Observe the character of the precipitate formed. 

Exp. 98. Fermentation of Glucose. — Fit a large acid bottle 
with a one-hole rubber stopper carrying a delivery tube. (See 
Fig. 123.) Connect this with a wash bottle containing lime 
water about two inches deep and let the delivery tube from the 
wash bottle dip into a beaker of lime water. Put into the acid 
bottle about a liter of water and about 200 c. c. of glucose syrup, 



Fig. 123. — Fekmentation of Glu- 
cose. The solution to be fer- 
mented is placed in the larger jar 
and is mixed with yeast. The 
smaller jar contains lime water 
to show that carbon dioxide as a 
product of the fermentation is es- 
caping through the delivery tube. 
In order that no carbon dioxide 
from the air may enter the lime 
water, a U-tube filled with so- 
dium hydroxide is joined to the 
exit tube. 



APPENDIX A 



315 



and to the mixture add a hundred cubic centimeters of brewers' 
yeast or a compressed yeast cake dissolved in a little water. In- 
sert the stopper into the mouth of the acid bottle and let the 
mixture stand in a warm room, observing- from time to time all 
changes that occur for several days. What do you learn from 
the lime water in the wash bottle? 
What is the use of the lime water 
in the beaker? Reserve the con- 
tents of the bottle for Exp. 100. 

Exp. 99. Effect of Heat Upon 
Sugar. — In a porcelain dish care- 
fully heat with continuous stirring 
a quantity of cane sugar until 
melted, and continue to increase 
the temperature to about 200° C. 
Carefully examine the product. 
How does it resemble sugar? How 
does it differ from sugar? 

Exp. 100. Distillation of Alcohol. 
— Arrange a distilling apparatus as 
shown in Figure 124 and put into 
the distilling flask about 500 c. c. 
of liquid decanted from the fermen- 
tation in Exp. 98. Continue the 
distillation till the thermometer 
reaches 99° C, then pour the con- 
tents of the receiver into a smaller 
distilling flask and distil again 
{see Fig. 125), continuing the pro- 
cess this time until only a small 
quantity of distillate has gone over, 
tillate on a glass rod into a flame; will it burn? Try a drop of 
alcohol in the same way. 

Exp. 101. Test for Alcohol. — In the solution to be tested for 
alcohol in a test tube or small flask dissolve about a gram of so- 
dium carbonate; heat the mixture nearly to boiling, and add 
small crystals of iodine, one at a time, so long as they will dis- 
solve to a colorless or nearly colorless liquid. Next let the solu- 




Fig. 124. — Distillation op Alco- 
hol FROM THE FERMENTEDMASH. 

The fermented material (or mash) 
is heated in the flask, F, in a 
water-bath. The distilling tube, 
G, is filled with broken glass or 
beads to condense the water and 
let it run back into the flask. The 
alcohol vapor is condensed in the 
condenser.C, which is but partly 
shown. The temperature is 
measured by the thermometer, T. 



Bring a drop of this dis- 



316 



A PRACTICAL CHEMISTRY 



tion cool. If alcohol was present in the solution, a yellow crystal- 
line precipitate of iodoform will appear. 

Exp. 102. Ethereal Salts. — Mix equal volumes of ethyl alco- 
hol and glacial acetic acid; then add a tenth of their volume of 
concentrated sulphuric acid and warm the mixture. Notice the 
odor of ethyl acetate which escapes. Repeat, using amyl alcohol. 
What does the odor resemble 1 ? 

Exp. 103. Fats and Oils in Food. — In a dry filter paper in a 
small funnel place some cornmeal; saturate the meal with ether, 




Fig. 125. — Redistillation of Alcohol. D is the distilling flask; T, the thermom- 
eter; R, the receiver; I and O are the inlet and outlet for the water which cools 
the condenser, C. 



and catch the ether in a dish as it filters through. (Note. — When 
using ether do not have any flames near.) Put the ether solution 
in the hood and let it evaporate. Examine what remains. Re- 
peat the experiment, using some walnut kernels or other nuts. 

Exp. 104. Nitrogenous Food or Albumen. — Put a little white 
of egg in a test tube and stand it in boiling water for a little 
time. What effect do you notice? Add concentrated nitric acid 
to the white of egg. Result ? Add a solution of mercuric chloride. 
Result? To a little white of egg add a little alcohol. Result? 

Exp. 105. Milk Preserving. — Put into each of three clean 



APPENDIX A 



317 



beakers 200 c. c. of sweet milk. Stand No. 1 in a pan of water 
and heat to 75° C, holding it at this temperature for about 20 min- 
utes, stirring from time to time. To No. 2 add a cubic centi- 
meter of formaldehyde, stirring meanwhile. Add nothing to 
No. 3. Cover all the beakers and let them stand in a warm place 
for three days, examining every 24 hours and tabulating results. 
What are the conclusions? 

Exp. 106. Coagulation of Milk. — In two test tubes put a little 
milk and add a few drops of hydrochloric acid to one and a little 
alcohol to the other and observe results. 

Into a clean beaker put 200 c. c. of sweet milk ; stand the beak- 
er in water and bring the temperature of the milk to about 35° C; 
dissolve half a rennet (junket) tablet in a few c. c. of water and 
stir the solution into the milk. Let the •» 

milk stand in a warm place for half an 
hour. Cut the curd thus formed with a 
knife and observe the separation of the 
whey. Taste each. 

Exp. 107. Testing Milk for Butter Fat. 
— Carefully weigh into a Babcock milk 
tester bottle (see Fig. 126) eighteen grams 
of milk; add 3 c. c. of a mixture of equal 
volumes of amyl alcohol and concentrated 
hydrochloric acid; then pour slowly into 
the bottle 10 c. c. of sulphuric acid (1.83 
specific gravity), shaking the bottle from 
time to time. Add enough of a freshly 
prepared mixture of equal volume of sul- 
phuric acid and water to bring the liquid 
up against the scale on the bottle neck. 
Balance the bottle in the machine and ro- 
tate at the rate of 1,000 revolutions per 
minute for four minutes. (See Fig. 127.) 
With a pair of dividers measure the length 
of the fat column from its bottom to the 

bottom of the meniscus. The scale is graduated to be read 
in per cent.; placing one leg of the dividers on the zero of the 




i^cy 



Fig. 126. — Babcock 
Milk Testing Bottle. 



318 



A PRACTICAL CHEMISTRY 



scale, measure off the length of the fat column and read in per 
cent, the fat contained in your sample of milk. 

Exp. 108. Vinegar. — Obtain some "mother of vinegar" from a 
sample of vinegar which has stood for some time and examine it 
under the microscope. Decolorize some vinegar by filtering it 
through animal charcoal. Neutralize with sodium hydroxide and 




Fig. 127. — Babcock Milk Tester. 



evaporate to crystallization. What are the crystals? Cool a 
bottle of 99 per cent, acetic acid to about 10° C. without shaking; 
admit a little air to the bottle and then shake it; observe the 
crystallization of the acetic acid. 

Exp. 109. Ferments in Bread. — Dissolve a yeast cake in a half 
cupful of water which has been heated to about 45° C. and stir 
in a cupful of flour. Let this dough stand at ordinary room 
temperature. Cut one-half of the dough into small pieces and 
place in a flask with an equal volume of water and heat nearly 
to the boiling point. Filter and test the filtrate for alcohol as 
described in Exp. 101. Let the other half of the dough stand 
another day and then examine it. Does it have any action on 
litmus? What is meant by "sour" bread? 

Exp. 110. Preparation of Sodium Bicarbonate. — Dilute 50 c. c. 
of strong ammonium hydroxide with twice its volume of water 



APPENDIX A 



319 



and shake this with a quantity of common salt till a cold satu- 
rated solution is obtained. Generate carbon dioxide, as in Exp. 
74, connect the delivery tube with a wash bottle, and slowly 
bubble the gas through the solution prepared above until a pre- 
cipitate of sodium bicarbonate has formed. {See Fig. 128.) 
Filter and dry this precipitate. What gas is formed when an 




Fig. 128. — Apparatus to Illustrate the S6lvay Process. In the Kipp apparatus 
carbon dioxide is generated by the action of hydrochloric acid in A upon the lime- 
stone in L. The carbon dioxide is washed by the water in the wash bottle, W, 
and then bubbles into the solution of salt in ammonium hydroxide in C. When 
the solution in C becomes saturated the sodium acid carbonate begins to pre- 
cipitate. 



acid is mixed with sodium bicarbonate? What effect does it have 
upon the bicarbonate to put it into hot water? (Try it.) 

Exp. 111. Baking Powder. — Thoroughly pulverize and mix 
two grams of cornstarch, five of sodium bicarbonate, and eleven 
of cream of tartar. Put water on a little of this powder in a 
beaker. Explain the action. What is the purpose of the corn- 
starch? What is the chemical name of cream of tartar? Make 



320 A PRACTICAL CHEMISTRY 

a dough of flour, water, and baking powder and let it stand for 
a few minutes. Explain result. 

Exp. 112. Examination of Baking Powders.— Examine several 
powders for starch by making them into paste and adding a few 
drops of iodine solution as in Exps. 95 and 96. To portions of 
the various powders add water, and shake till the gas ceases to 
come off and then filter. Test portions of the filtrates for sul- 
phates and phosphates as described in previous experiments. 
Test a third portion for tartrates as follows: Make a 10 per 
cent, solution of silver nitrate, and to this add ammonium hydrox- 
ide, drop by drop, till a brown precipitate of silver oxide forms. 
Pour this mixture into a test tube to the depth of one inch and 
add an equal volume of the solution to be tested. Heat the mix- 
ture to boiling. If a silver mirror forms on the tube, it may be 
taken as evidence that a tartrate or tartaric acid was present in 
the baking powder. On charcoal oxidize a sample of each pow- 
der which does not contain a phosphate with the oxidizing flame, 
and then moisten the residue with a drop of a solution of cobalt 
nitrate. Again heat the mass with the oxidizing flame. If the 
substance turns blue, aluminum is present, if phosphates are 
absent. The presence of a sulphate of aluminium indicates that 
the powder contained an alum. Tabulate your results, indicating 
which powders are alum, which phosphate, and which tartrate 
powders. 

Exp. 113. Formaldehyde.— Fit a large test tube with a cork 
stopper and put into the tube a few drops of methyl alcohol. 
Make a piece of copper wire gauze into a roll, and heat it to 
a high temperature in the blast lamp. .Plunge the hot copper 
gauze into the test tube; carefully notice the odor of the vapor 
which escapes from the tube, being careful not to inhale too 
much of the vapor. 

Exp. 114. Antidote for Mercuric Chloride. — Shake up a little 
white of egg in a test tube with a little of a solution of corrosive 
sublimate, mercuric chloride, and observe the result. Explain how 
the white of egg serves as an antidote. 

Exp. 115. Iodoform. — Put into a small flask 10 c. c. of alcohol 
and three times its volume of water. In this dissolve five grams 
sodium carbonate and bring the solution to boiling; then add 



APPENDIX A 321 

iodine, crystal by crystal, so long as it will dissolve. Let the 
solution cool, filter out the precipitate and recrystallize it from 
alcohol. Notice the odor of the iodoform. 

Exp. 116. Borax Beads. — Make a round loop on the end of a 
platinum wire by bending it around the lead of a pencil till it 
forms an open circle; heat this in the Bunsen flame and dip it 
into some borax. Heat the borax which adheres to the wire till 
it forms a transparent glass bead. Put the smallest possible 
quantity of a cobalt salt upon this bead and heat again till trans- 
parent. If the bead is not colored, add more of the salt and 
repeat the heating. Observe the color of the bead. Shake this 
bead off while melted, and clean the wire by making new beads 
and shaking them off till a colorless bead is obtained. Repeat, 
adding a very small quantity of a chromium salt to the bead, 
and observe the color. 

Exp. 117. Boric Acid Crystals. — Make a hot saturated solu- 
tion of borax and add to it one-fourth its volume of sulphuric 
acid. Let the solution cool and filter out the crystals that form. 
What are these crystals? 

Exp. 118. Test for Boric Acid. — Mix some boric acid with al- 
cohol and ignite the alcohol. Notice the green color of the flame. 
Can you get the same test with borax? Mix the borax with sul- 
phuric acid and then make the test. 

Exp. 119. Making Chloroform. — Arrange apparatus as for dis- 
tilling alcohol. Put into the flask 50 c. c. of alcohol, 200 c. c. of 
water, 100 grams of bleaching powder, and 50 grams of lime. 
Distil the mixture. Examine the heavy, oily liquid which collects 
at the bottom of the receiver and notice its odor. Shake up in 
a test tube a little sodium hydroxide solution with a few crystals 
of chloral hydrate. Is the odor that of chloroform? 

Exp. 120. Making Ether. — In an evaporating dish put 100 
c. c. of alcohol; place the dish in a pan of cold water and slowly 
stir into it one-half its volume of concentrated sulphuric acid. 
Arrange apparatus (see Fig. 129) as for distilling, but pass a 
dropping funnel through the stopper in the distilling flask. Pour 
the mixture of acid and alcohol into the distilling flask and alcohol 
into the dropping funnel. Heat the acid mixture to about 140° 
C; then slowly let the alcohol from the dropping funnel fall 



322 



A PRACTICAL CHEMISTRY 



into the flask. When the distillation is completed add to the dis- 
tillate an equal volume of water, and separate the layers of 
water and ether by means of a separatory funnel. Notice the 
odor of the ether. Remember the caution about keeping flames 
at a distance from ether. Why this precaution? What are the 
properties of ether? 




Fig. 129. — Making Ether. The mixture of alcohol and sulphuric acid is put into 
the flask D, and heated till the thermometer shows a temperature of 140° C. 
Alcohol is then allowed to slowly drip from A while the heating is continued. 
Ether distils over and collects in E. The screen, S, is to prevent ether vapor 
from reaching the flame, F. C, the condenser, is kept cool by water flowing in 
at I and out at O. 



Exp. 121. Preparation of Iodine. — Pulverize and mix equal 
parts of potassium iodide and manganese dioxide, and place the 
mixture in a test tube without letting it get onto the walls of the 
tube. Through a thistle tube add enough sulphuric acid to cover 
the mixture. Carefully heat till the upper part of the tube is 
coated with crystals of iodine. 

Exp. 122. Re crystallizing Iodine. — Scrape out a quantity of 
the iodine crystals and place in a clean, dry test tube; heat care- 
fully; observe the color of the vapor and the character of the 
crystals which form upon the sides of the test tube. 

Exp. 123. Tincture of Iodine. — Dissolve some iodine in alco- 



APPENDIX A 



323 



£ 



% 




hoi; paint some of this tincture upon your hand and notice the 
color. (This stain may be removed by ammonium hydroxide.) 
Is iodine soluble in water? Try its solubility in a solution of 
potassium iodide. 

Exp. 124. Preparation of Bromine. — Fit a test tube with a 
delivery tube bent twice at right angles. Pulverize and mix equal 
parts of potassium bromide and manganese dioxide, and put the 
mixture into the test tube. Add enough sulphuric acid to cover 
the mixture, insert the stopper, and let the delivery tube extend 
into a test tube containing about an inch 
of water. The delivery tube should near- 
ly reach the water, but not dip into it. 
{See Fig. 130.) Heat the test tube con- 
taining the mixture carefully; observe 
the bromine vapor as it ascends the 
tube; also the bromine at the bottom of 
the water and the bromine water above. 
(Caution: Do not let the bromine get 
upon your skin.) 

Exp. 125. Calomel. — Mix solutions of 
mereurous nitrate and salt together and 
filter out the precipitate formed. Try 
the solubility of this precipitate in hot 
water and in ammonium hydroxide. 
How can you tell calomel from corrosive 
sublimate? Do not throw mercury salts 
in the sinks. What antidote would you 
give for mereurous nitrate? 

Exp. 126. Action of Acids and Alkalies on Fabrics. — Place 
strips of cotton and wool in solutions of acids and bases, and 
boil for a minute. Carefully observe results. Can you devise a 
method of distinguishing between cotton and wool? 

Exp. 127. Dyeing. — Boil a gram of logwood extract with 200 
c. c. of water till dissolved. In an evaporating dish boil a bit 
of cotton cloth with some of the above solution of logwood for 
a minute. Remove the cloth, wash and dry, and keep it for 
comparison. 

Make a lake by adding a few drops of the dye to a concen- 



Fig. 130. — Liberating 
Bromine. The test tube, 
K, contains the mixture 
of manganese dioxide, po- 
tassium bromide and sul- 
phuric acid. The deliv- 
ery tube, T, is of glass 
and extends to within a 
half inch of the water. 
The liquid bromine col- 
lects at B. 



324 A PRACTICAL CHEMISTRY 

trated solution of alum, and then add ammonium hydroxide till 
the solution is alkaline. The precipitate is the lake. Filter it out. 
What does it contain 1 ? 

Dye a second bit of cotton goods, using solutions of ferric 
chloride and ammonium hydroxide to produce a mordant. Wet 
the goods with the solution of ferric chloride and wring out the 
excess of liquid; next saturate it with ammonium hydroxide and 
again squeeze out the excess. Next boil the goods with the dye; 
then wash and dry it. Compare the result with that obtained by 
dyeing without the mordant. 

Compare the action of a dye on cotton and wool thus. Make 
a saturated solution of picric acid, and in this boil for a few 
minutes a bit of cotton goods and some strands of woolen yarn. 
After removing them wash them thoroughly in cold water. What 
difference do you find? What is the use of a mordant? 

Exp. 128. Solvents for Cleaning. — Stain five bits of cotton 
goods with ink, five with fruit, five with blood, five with paint, 
and five with iron rust. Wash one sample of each set with cold 
water, one with hot water, one with gasoline or carbon tetrachlo- 
ride, one with a concentrated solution of oxalic acid, and one 
with dilute hydrochloric acid, and record results in tabular 
form. 

Exp. 129. Soap Making. — In an evaporating dish boil a mix- 
ture of five grams of cocoanut oil, ten grams of sodium hydrox- 
ide, and 75 c. c. of water. As the water boils away add more 
boiling water, and continue the process till all of the oil seems to 
be dissolved. With continuous stirring add ten grams of common 
salt, and set the dish aside to cool. When cold collect the soap 
from the top of the liquid and dissolve it in a small beaker of 
water, and save the solution for Exp. 130. 

Exp. 130. Action of Acids and Salts Upon Soap. — To a portion 
of the soap solution made in Exp. 129 add a few drops of hydro- 
chloric acid; carefully examine the precipitate of solid acids. 
Explain the formation of these acids. To a second portion add 
a solution of calcium chloride, to another portion a solution of 
copper sulphate, and to a third some sodium or potassium sul- 
phate. Explain the results. 

Exp. 131. Hard and Soft Water. — Make a soap solution by 



APPENDIX A 325 

dissolving ten grams of good white soap in 50 c. c. of alcohol, 
and then add an equal volume of water. Make permanent hard 
water by dissolving ten grams of calcium chloride in one liter of 
water. Make temporary hard water by diluting lime water 
with an equal volume of distilled water, and then adding car- 
bonated water in such quantity that the precipitate which first 
forms dissolves. 

To 100 c. c. of distilled water add the soap solution, drop by 
drop, with violent shaking to determine how many drops are re- 
quired to form a lather that will cover the whole surface of the 
liquid. Repeat, using both kinds of hard water. Boil 100 c. c. 
of the temporary hard water for two minutes, and then test with 
the soap solution. Does it take as much soap as before boiling 1 ? 
How does the quantity compare with that required by distilled 
water? Why is temporary hard water so named 1 ? To 100 c. c. 
of permanent hard water add two grams of washing soda and 
shake till dissolved; then test with soap solution as before. Re- 
sult? What are the disadvantages of hard water? In how many 
ways can hard water be softened? Test your drinking water for 
temporary and permanent hardness. 

Exp. 132. Hydrofluoric Acid. — In a leaden dish put a spoonful 
of fluorspar and enough sulphuric acid to make a paste. Coat a 
glass plate with wax and engrave in the wax a design ; place the 
glass with the wax down over the dish and let stand for an hour 
or more. (Note. — In working with hydrofluoric acid care should be 
taken not to breathe the fumes or get the solution upon the skin.) 
Glass is best coated with wax by dipping it into the melted wax. 
The wax may be removed by putting the glass into hot water, 
and when the wax has melted pouring water and wax off together. 
Remove the wax from the glass used in this experiment. Explain 
the action which has produced the result. Glass may be etched 
by painting the design in the wax over with a solution of hydro- 
fluoric acid in water in place of the vapor as used in this experi- 
ment. 

Exp. 133. Water Glass. — Mix equal weights of finely pow- 
dered sand, sodium carbonate, and potassium carbonate and put 
them into a sand crucible and heat to gentle fusion in a small 
furnace until the gas bubbles have nearly stopped coming off. 



326 A PRACTICAL CHEMISTRY 

What products are formed in this reaction? Pour the water 
glass into about four times its volume of water, thus dissolv- 
ing it to a thick liquid. Dilute some water glass and pour 
into it some hydrochloric acid. What is the precipitate which 
forms'? 

Put a little water glass between two pieces of glass and press 
them together; then let them stand till dry. What use of water 
glass does this suggest? Make other silicates by mixing solutions 
of other salts, such as copper sulphate, silver nitrate, and man- 
ganese sulphate, with solutions of water glass ; examine and name 
the silicate formed in each case. Make a diluted solution of 
water glass in a beaker and drop in crystals of copper sulphate, 
nickel nitrate and zinc sulphate, and let stand for an hour. 
What is osmotic pressure? How does it enter into this experi- 
ment? 

Exp. 134. Quicklime. — Test a piece of limestone with wet 
litmus paper; then heat one end of it in the oxidizing blowpipe 
flame or blast lamp for some minutes and test again with litmus. 
Read text and explain the change which has occurred. 

Exp. 135. Slaked Lime. — Add hot water, drop by drop, to a 
lump of quicklime in an evaporating dish. Why does the lump 
grow larger? What becomes of the water? How can you make 
limewater and milk of lime? 

Exp. 136. Mortar. — Mix lime and water to a thin paste, and 
then add clean sand until a stiff, mushlike paste is obtained. Set 
this mortar aside for a week or more to harden and then test it 
as follows for carbonates. Put a little limewater into one test tube 
and some of the mortar to be tested into the other. Add a little 
dilute hydrochloric acid to the mortar and decant into the lime- 
water any gas which may escape. {See Fig. 131.) Close the 
limewater tube with your thumb and shake it. What evidence 
do you have of carbonates? Why should the mortar contain a 
carbonate ? 

Exp. 137. Plaster of Paris. — Carefully heat some crystals of 
gypsum in a test tube, holding the mouth of the tube inclined 
downward. What two compounds are thus obtained? To some 
plaster of Paris add enough water to make it plastic and mold it 
into a cast. Let it stand till hard. How do you explain its, 



APPENDIX A 



327 




hardening? How do you explain the setting of cement? What 
experiment have we had before which was related to this one? 

Exp. 138. Tests for Iron in Ferrous and Ferric Compounds. 
— Make a solution of ferrous sulphate by dissolving the crystals 
in boiling water. Also make a solution of ferric chloride. Put 
about 2 c. c. of the ferrous solution in each of three test tubes. 
To one add a few drops of potassium 
ferrocyanide, to another a little of a 
solution of potassium ferricyanide, 
and to the third a few drops of a 
solution of ammonium sulphocyanate. 
Tabulate the results. Repeat, but 
this time use the ferric chloride in 
place of the ferrous sulphate. Care- 
fully examine the tabulated results 
and choose those which are character- 
istic for tests for ferrous and ferric 
compounds. 

Exp. 139. Change of Valence of 
Iron in Compounds. — Iron in ferrous 
compounds has a valence of two and 

in ferric a valence of three. To a solution of ferrous sulphate 
add a little nitric acid and boil for three minutes. Test the 
solution for ferrous and ferric iron. To a solution of ferric 
chloride add hydrochloric acid and metallic iron, and heat nearly 
to boiling for ten minutes; then test the solution for ferrous and 
ferric iron. .How is valence increased? How made less? 

Exp. 140. Hydroxides of Iron. — To a solution of ferrous sul- 
phate add ammonium hydroxide until the solution smells strongly 
of it. Examine the precipitate of ferrous hydroxide. Close the 
tube with the thumb and shake up the contents, admitting air 
from time to time and noticing all changes. In another tube add 
ammonium hydroxide to a solution of ferric chloride and examine 
the ferric hydroxide. What was formed by shaking the ferrous 
hydroxide with air? 

Exp. 141. Stannous and Stannic Chloride. — Put some granu- 
lated tin into hydrochloric acid. What gas is given off? The 
solution contains stannous chloride; how can you convert it into 



Fig. 131. — Test for Carbo- 
nates. The substance to be 
tested is put into the tube, C, 
with a mixture of equal vol- 
umes of water and hydro- 
chloric acid while limewater 
is put into L. The heavy 
carbon dioxide is poured into 
L, which is then closed with 
the thumb and shaken. 



328 A PRACTICAL CHEMISTRY 

stannic chloride? Mix stannous chloride with mercuric chloride, 
being careful to have the mercuric chloride in large excess. How 
can you prove that the white precipitate is calomel? How do 
you explain its formation? Add the stannous chloride in excess 
and warm the mixture. Collect the precipitate and rub it upon 
bright copper. What is it? Where did it come from? 

Exp. 142. Aluminium as a Reducing Agent. — Make a mixture 
of ferric oxide and powdered aluminium in the proportion of 
one of aluminium to three of ferric oxide, and put the mixture 
into a sand crucible. At the top, in the center of the crucible, 
place a small quantity of powdered magnesium mixed with a 
small quantity of potassium chlorate. With a long taper light 
this magnesium mixture, taking care to keep the face away from 
the crucible. When the crucible has cooled crack it open and 
examine the contents. How do you explain this reaction? 

Exp. 143. Fusible Metal. — Mix in a crucible 8 parts of lead 
with 5 parts of bismuth, 4 parts of tin and 3 parts of cadmium, 
and carefully fuse. After the fused mass has become thoroughly 
mixed let the alloy cool, and when hard place in a beaker of 
water, stir with a thermometer and carefully heat. Note the tem- 
perature at which fusion takes place. What uses can you suggest 
for an alloy melting at this temperature? 

Exp. 144. Amalgams. — Amalgamate several metals, as zinc, 
copper, and silver, by rubbing them (being certain that they are 
first clean) with mercury, or by placing them in solutions of salts 
of mercury. 

Exp. 145. Test for Gold. — Dissolve the substance to be tested 
in aqua regia. Make a mixture of equal quantities of stannous 
and stannic chlorides, and add a few drops of the solution to be 
tested. Stand it aside for a short time and observe the purple 
precipitate of finely divided gold. 

Exp. 146. Silver Salts. — Dissolve a dime in nitric acid. What 
salt of silver is formed? Why is the solution blue? What 
colored gas is given off? Dilute the solution and add a solution 
of common salt so long as a precipitate forms. What is this 
precipitate? Save it for the next experiment. 

Exp. 147. Silver Plating. — Thoroughly wash the precipitate 
of silver chloride made in the last experiment. Add a ten per 



APPENDIX A 329 

cent, solution of potassium cyanide until the silver chloride is com- 
pletely dissolved, and then add an additional 50 c. c. of the cyan- 
ide solution. (Caution: Remember that potassium cyanide is a 
dreadful poison.) Thoroughly clean the object to be plated and 
make it the cathode, using a dry cell as the source of current. 
Make a plate of pure silver the anode. Close the circuit and 
allow the action to continue until a white deposit begins to show 
on the object being plated. Remove the object and polish till 
bright. Repeat this process several times till a plating of desired 
thickness is obtained. 



APPENDIX B 



CHEMICAL ARITHMETIC 

In most practical chemical work calculations must be made to 
determine the quantities of substances involved. In the following 
pages are illustrations of most of the various sorts of calculation 
required. 

I. THE METEIC SYSTEM 

In all scientific work the metric system of weights and meas- 
ures is used. This may be found discussed in any arithmetic, but 
for convenience the tables of length, capacity, and weight are 
given: 



Table of Linear 

Myriameter = 

Kilometer (km.) = 

Hectometer = 

Dexameter = 
Meter (m.) 

Decimeter (dm.) = 

Centimeter (cm.) = 

Millimeter (mm.) = 

Mikron (/*.) = 



Measure (Length) 

10,000 meters 

1,000 meters 

100 meters 

10 meters 

0.1 of a meter 
0.01 of a meter 
0.001 of a meter 
0.000001 of a meter 



Table of Cubic Measure (Capacity) 



Cubic hektometer = 

Cubic dekameter = 
Cubic meter (m\) 

Cubic decimeter (dm. 3 ) = 
Cubic centimeter (cm. 3 or c. c.) = 

Cubic millimeter (mm. 3 ) = 
Liter (1.) 



1,000,000 cubic meters 
1,000 cubic meters 

0.001 of a cubic meter 
0.000001 of a cubic meter 
0.000000001 of a cubic meter 
1,000 cubic centimeters 



330 



APPENDIX B 


Table 


OP 


Weights 


Metric ton (t.) 


= 


1,000,000 grams 


Quintal (q.) 


= 


100,000 grams 


Miriagram 


= 


10,000 grams 


Kilogram (kg.) 


= 


1,000 grams 


Hektogram 


= 


100 grams 


Dekagram 


= 


10 grams 


Gram (g.) 






Decigram (dg.) 


= 


0.1 of a gram 


Centigram (eg.) 


= 


0.01 of a gram 


Milligram (mg.) 


= 


0.001 of a gram 


Mikrogram (y) 


= 


0.000001 of a gram 


Approximate Equivalents 


Meter 


= 


39.37 inches = 3^ feet 


Kilometer 


= 


3/5 of a mile 


Liter 


= 


1 quart 


Kilogram 


= 


2 1/5 pounds, avoirdupois 


Gram 


= 


153^ grains 


Hectare 


= 


2)4, acres 


Square Meter 


= 


10 square feet 



331 



It will be observed that the metric system is a decimal system. 



II. MOLECULAE WEIGHTS 



From Dalton's atomic theory we learned that each element is 
made up of atoms and that the atoms of the various elements 
have different weights. It must follow that the atoms of each 
element have definite weight. It is also evident that the weights 
of these atoms cannot easily be determined in grams or pounds. 
Their weights, however, as compared with some standard atom, 
can be definitely found. The weight of the hydrogen atom 
and the one-sixteenth of the weight of an oxygen atom have been 
chosen as standard units with which to compare the weights of 
all other atoms. These two units are not quite of equal value; 
when the oxygen atom is valued at 16, the atomic weight of 
hydrogen is 1.008. At present most chemists prefer the oxygen 
standard. 



332 



A PRACTICAL CHEMISTRY 



List op the Elements, Their 

0: 

Aluminium Al 27.1 

Antimony .Sb 120.2 

Argon A 39.88 

Arsenic As 74.96 

Barium Ba 137.37 

Beryllium Be 9.1 

Bismuth Bi 208.0 

Boron B 11.0 

Bromine Br 79.92 

Cadmium Cd 112.4 

Caesium Cs 132.81 

Calcium Ca 40.07 

Carbon C 12.0 

Cerium Ce 140.25 

Chlorine CI 35.46 

Chromium Cr 52.0 

Cobalt ...Co 58.97 

Columbium Cb 93.5 

Copper Cu 63.57 

Dysprosium . . . .Dy 162.5 

Erbium Er 167.7 

Europium Eu 152.0 

Fluorine F 19.0 

Gadolinium Gd 157.3 

Gallium Ga 69.9 

Germanium Ge 72.5 

Gold Au .197.2 

Helium He 3.99 

Holmium ...... Ho 163.5 

Hydrogen H 1.008 

Indium In 114.8 

Iodine I 126.92 

Iridium Ir 193.1 

Iron Fe 55.84 

Krypton Kr 82.92 

Lanthanum La 138.0 



Symbols, and Atomic Weights 

= 16 

Lead Pb 207.10 

Lithium Li 6.94 

Lutecium Lu 174.0 

Magnesium Mg 24.32 

Manganese Mn 54.93 

Mercury Hg 200.6 

Molybdenum ... .Mo 96.0 

Neodymium Nd 144.3 

Neon Ne 20.2 

Nickel Ni 58.68 

Nitrogen N 14.01 

Osmium Os 190.9 

Oxygen O 16.0 

Palladium Pd 106.7 

Phosphorus P 31.04 

Platinum Pt 195.2 

Potassium K 39.10 

Praseodymium . . Pr 140.6 

Radium Ra 226.4 

Rhodium Rh 102.9 

Rubidium Rb 85.45 

Ruthenium Ru 101.7 

Samarium Sm 150.4 

Scandium Sc 44.1 

Selenium Se 79.2 

Silicon Si 28.3 

Silver Ag 107.88 

Sodium Na 23.0 

Strontium Sr 87.63 

Sulphur S 32.07 

Tantalum Ta 183.5 

Tellurium Te 127.5 

Terbium .Tb 159.2 

Thallium Tl 204.0 

Thorium Th 232.4 

Thulium Tm 168.5 



APPENDIX B 



333 



Tin Sn 119.0 

Titanium Ti 48.1 

Tungsten W 184.0 

Uranium U 238.5 

Vanadium V 51.0 



Xenon Xe 130.2 

Ytterbium Yb 172.0 

Yttrium Yt 89.0 

Zinc Zn 65.37 

Zirconium Zr 90.6 



The calculation of the molecular weight when the formula of 
the compound is known is an easy matter, and consists only in 
adding the weights of the various atoms found in the molecule. 
Thus, if it is required to find the molecular weight of the com-" 
pound having the formula, C 6 H 10 O 6 , by reference to the table 
we find the atomic weight of carbon = 12, hydrogen = 1.008 
(which for our purposes we may call 1), and oxygen = 16. 

Then 

C 6 = 12X 6 = 72 

Hio = IX 10 = 10 

6 = 16X 6 = 96 

Molecular weight of C 6 H 10 O 6 = 178 

Problems 

1. Calculate the molecular weight of benzene, C 6 H 6 ; of cane 
sugar, C^H^O^; of zinc sulphate, ZnS0 4 . 

2. What is the molecular weight of copper sulphate, CuS0 4 . 
5H 2 0? 

3. Find the molecular weight of Pb(N0 3 ) 9 and of lead acetate^ 
Pb(C 2 H 3 2 ) 2 .3H 2 0. 

III. PEKCENTAGE COMPOSITION 

The percentage composition of compounds is usually obtained 
as a result of analysis, but sometimes it is desirable to know the 
percentage of an element in a given compound when the formula 
is known. For example, what percentage of oxygen is contained 
in potassium chlorate 1 ? First we find the molecular weight of 
potassium chlorate, KC10 3 , as shown above. 



334 



PRACTICAL 


CHEMISTRY 


K 


= 


39.10 


CI 


= 


35.46 


3 


= 


48. 



KC10, = 122.56 

It will be observed that, while the total molecular weight of the 
compound is 122.6, 48 of this is due to the oxygen, or, in other 

words, the oxygen is — — — of the compound. Reducing this to 

1226 
hundredths or per cent., we have 480.00 -r- 1226 = 39.15 per cent. 
of oxygen. The percentage of the other elements is found in the 
same way. Thus, to find the percentage of potassium in this 

391 
compound we would have potassium = of the molecular 

weight of the compound, and, reducing this to hundredths or per 
cent., 3,915 -f- 12,260 = 31.93 per cent, potassium. 

Problems 

1. What percentage of lead is found in red lead, Pb 3 4 ? 

2. Calculate the percentage composition of sodium carbonate, 
Na 2 C0 3 ; of oxalic acid, C 2 H 2 4 . 

3. What weight of hydrogen may be got from 100 grams of 
hydrochloric acid, HCl? 

4. How does the percentage of oxygen in nitrous oxide com- 
pare with the percentage in nitrogen peroxide? 

5. How much sulphur would be required to make one ton of 
ferrous sulphide? 

IV. TO CALCULATE THE QUANTITY OF A GIVEN ELE- 
MENT OR COMPOUND REQUIRED IN A CHEM- 
ICAL REACTION 

From the previous section it may be seen that it is possible to 
calculate the quantity of a given element or compound required in 
a given reaction. The work may be somewhat simplified by the 
use of proportion. 

The first step in working problems of this sort is to ascertain 
the chemical equation which represents the reaction. The molec- 



APPENDIX B 335 

ular weights of the substances involved is next calculated, and 
finally the proportion is formed and solved. For example, it is 
required to know what weight of zinc reacting with sulphuric acid 
will give 100 grams of hydrogen. First the chemical equation is : 

Zn + H 2 S0 4 = ZnS0 4 + H 2 
Although several chemical substances are involved in this reac- 
tion, the only things which need to be considered in the calcula- 
tion are the zinc and hydrogen, since it will be observed that one 
atom of zinc liberates two atoms of hydrogen. From the table 
we see that a zinc atom weighs 65.4, while two atoms of hydrogen 
weigh 2. Therefore, 65.4 grams of zinc produce 2 grams of 
hydrogen and x grams of zinc produce 100 grams of hydrogen, or, 

65.4 2 

x " 100 
Solving for x, 

2x = 6540 

x = 3270 grams of zinc 
If the problem had been, "What weight of hydrogen can be gen- 
erated by the action of 1,000 pounds of zinc on sulphuric acid?" 
the chemical reaction would have been the same as above, and 
we would have made the statement : 65.4 pounds of zinc produce 
2 pounds of hydrogen, and 1,000 pounds of zinc will generate 
x pounds of hydrogen, or, 

65.4 2 

1000 ~ x 
Solving for x, 

65. 4x = 2000 

x = 30.5 + lbs. 
In solving the problem "What weight of manganese dioxide will 
be required to liberate 100 grams of iodine from potassium 
iodide?" we have the equation: 

2KI +Mn0 2 + 2H 2 S0 4 = 21 + MnS0 4 + K 2 S0 4 + 2H 2 

We require only the Mn0 2 and 21 for our calculation. The 
molecular weight of Mn0 2 = (Mn = 55 + 2 — 32) =87, and 
21 = 2 X 127 = 254; hence 87 grams of Mn0 2 will liberate 254 



336 A PRACTICAL CHEMISTRY 

grams of iodine, and x grams of Mn0 2 will liberate 100 grams 
of iodine, or, 

_87 254 

x " 100 
254x = 8700 

x = 34 -f- grams 

If the problem were, "(a) What weight of manganese dioxide is 
required to react with 100 grams of potassium iodide to liberate 
iodine, and (b) what weight of potassium sulphate may be ob- 
tained as a by-product?" the solution would be similar to the 
above. The same chemical reaction is involved and may be rep- 
resented by the same chemical equation. In (a) we find that 
2KI and Mn0 2 are the molecules involved. 

The weight of 2KI = 2K = (2X39.15) = 78.3 
21 = (2X127) = 254 
2KI = 332.3 

The weight of Mn0 2 (found above) = 87 

87 grams of Mn0 2 react with 332 .3 grams of KI, 
therefore, x grams of Mn0 2 react with 100 grams 
of KI, or 

87 332.3 

x " 100 
Which may be solved as above for x. 

To solve (b) we must know the molecular weights of potassium 
sulphate, K 2 S0 4 , and of two molecules of potassium iodide, 2KI. 

K 2 = 78.3 
S = 32. 

Q 4 = 64. 2KI (as found above) = 332 . 3 

K 2 S0 4 = 174.3 
Then for every 332.3 grams of potassium iodide entering into the 
reaction 174.3 grams of potassium sulphate are formed, and for 
100 grams of potassium iodide x grams of potassium sulphate 
will be formed, or, 

332J5 174.3 

100 x 

This is solved for x as in previous examples. 



APPENDIX B 337 



Problems 



1. (a) How many pounds of sodium reacting with water 
would be required to generate 100 pounds of hydrogen ? (b ) What 
weight of sodium hydroxide will be formed at the same time? 

2. (a) Calculate what weight of sodium sulphate, Na 2 S0 4 , 
may be made from 1,000 grams of sodium hydroxide 1 ? (b) How 
many grams of water would be formed in the reaction? 

3. What weight of sulphur is required to produce one ton of 
sulphur trioxide? To make one ton of sulphuric acid? 

4. How does the weight of hydrogen which can be produced 
by the action of 100 ounces of sodium upon water compare with 
the weight of hydrogen produced by the action of 100 ounces of 
potassium upon water? 

5. How much more will it cost to generate 100 pounds of 
hydrogen by the action of zinc on hydrochloric acid than by the 
action of iron on hydrochloric acid, when iron costs two cents 
per pound and zinc nine cents per pound? 

6. What weight of potassium permanganate will be needed to 
oxidize one thousand pounds of hydrochloric acid according to 
the equation, 

2KMn0 4 + 16HC1 = 2MnCl 2 + 2KC1 + 8H 2 + 5C1 2 ? 

7. (a) What weight of manganous chloride would be formed 
during the production of 1,000 pounds of chlorine? (b) What 
weight of potassium chloride? 

8. What weight of sulphuric acid is required to react with 
one ton of salt in the manufacture of hydrochloric acid? 

9. What weight of sulphuric acid will react with 100 pounds 
of sodium hydroxide? 

10. If one hundred pounds of sodium hydroxide are dissolved 
in one hundred gallons of water, what weight of sulphuric acid 
will be required to neutralize one-half of it? 

11. If 100 pounds of sulphuric acid be dissolved in ten gal- 
lons of water, what portion of it will be required to neutralize 
one gallon of the solution of sodium hydroxide mentioned in 
question 10? 

12. If 40 grams of sodium hydroxide be dissolved in 1,000 



338 A PRACTICAL CHEMISTRY 

c. c. of water, and 36.45 grams of hydrochloric acid be dissolved 
in 1,000 c. c. of water, how much of the sodium hydroxide solution 
will be neutralized by 1 c. c. of the hydrochloric acid ? 

13. If 10 c. c. of the above solution of sodium hydroxide will 
neutralize 5 c. c. of another solution of hydrochloric acid, what 
weight of acid is thus neutralized? 

14. If 25 c. c. of the solution of sodium hydroxide mentioned 
in question 12 can be neutralized by 50 c. c. of a solution of 
sulphuric acid, what weight of sulphuric acid is in one liter of 
this acid? 

15. If 100 grams of sulphuric acid be diluted to one liter, and 
33 c. c. of this dilute acid will neutralize 99 c. c. of a solution of 
potassium hydroxide, what is the concentration of this solution 
of potassium hydroxide? 

16. If a solution of sodium hydroxide contains 20 grams of 
sodium hydroxide per liter, how many c. c. of this solution will 
be required to neutralize one liter of a solution of hydrochloric 
acid containing 10 grams per liter? 

17. What is the concentration of a solution of calcium hydrox- 
ide, 1 c. c. of which will neutralize 1 c. c. of a solution of acetic 
acid having a concentration of 50 grams per liter? 

V. VOLUME OF GAS OBTAINED FEOM GIVEN WEIGHT 
OF CHEMICALS 

When the volume of gas is required it may be obtained by 
dividing the weight of gas by the weight of one liter of the gas. 
Since the weights of gases vary with temperature and pressure, 
their weights are usually stated as measured under standard con- 
ditions, i. e., at 0° C. and under a pressure of 760 mm. of mer- 
cury. The weight of one liter of some of the more common 
gases measured under standard conditions is given in the following 
table: 

Acetylene 1 . 1614 grams 

Air 1 .2923 grams 

Ammonia 0.7617 gram 

Carbon dioxide 1 . 9641 grams 

Carbon monoxide 1 . 2499 grams 

Chlorine 3 . 1650 grams 



APPENDIX B 339 

Hydrochloric acid 1 . 6275 grams 

Hydrogen. . . 0.08984 gram 

Methane . 7157 gram 

Nitric oxide 1 . 3410 grams 

Nitrogen 1 . 2501 grams 

Oxygen 1 .4285 grams 

Problems 

1. What volume of hydrogen measured under standard con- 
ditions can be obtained by the action of zinc upon 1,000 grams 
of hydrochloric acid? 

2. Which will give the greater volume of oxygen, 200 grams 
of mercuric oxide or 200 grams of potassium chlorate? 

3. (a) What volume of chlorine can be obtained from the 
decomposition of 1,000 grams of salt by electrolysis? (b) What 
volume of hydrogen will be obtained at the same time? 

4. What weight of ammonium nitrite will be required to gen- 
erate 100 liters of nitrogen? 

5. What volume of oxygen will be required to burn com- 
pletely 100 liters of acetylene? 

6. How much greater volume will be occupied by 100 grams 
of ammonia than by 100 grams of chlorine, both gases being 
measured under standard conditions? 

VI. TEMPERATURE CHANGES 

It sometimes becomes necessary to change temperature readings 
as given by the centigrade thermometer into the corresponding 
degrees of the Fahrenheit thermometer, and vice versa. 

By a comparison of the scales of the two thermometers it will 
be observed that, while on the centigrade scale there are 100 de- 
grees between the freezing and boiling points, on the Fahrenheit 
the same space is divided into 180 degrees. Hence, 100 centigrade 
degrees are equal to 180 Fahrenheit, or, 

5°C. = 9° F. 

1°C. = 9/5°F. 

and 9°F. = 5° C. 

1°F. = 5/9°C. 



340 A PRACTICAL CHEMISTRY 

Since 32° F. corresponds to 0° 0.', it is necessary to subtract 
32° from the Fahrenheit reading before multiplying by 5/9 to 
obtain the corresponding centigrade reading. In like manner to 
obtain the Fahrenheit reading corresponding to a given centi- 
grade temperature it is necessary to multiply the number of 
centigrade degrees by 9/5 and add 32° to the result. 

Illustrations. — What temperature centigrade corresponds to 68° 
Fahrenheit ? 

5/9 of (68 — 32) = 5/9 of 36 = 20° C. 

Change 20° C. to the corresponding Fahrenheit temperature. 

Since 1° C. = 9/5° F., 

20° C. = 9/5 of 20 = 36 Fahrenheit degrees, 

but since the zero on the Fahrenheit scale is 32° below the centi- 
grade zero, 32° must be added to the 36 obtained in the above 
calculation, making the Fahrenheit temperature 68°. 

When the temperature is below zero the subtraction and ad- 
dition of the 32 degrees must be done algebraically. Thus, 
change — 123° F. to the corresponding centigrade temperature. 
Subtracting 32 from — 123 we have — 155, and 

5/9 of —155 = - 86 1/9° C. 

Likewise to change — 200 C. to the corresponding Fahrenheit 
temperature : 

9/5 of —200 = —360 

and —360 + 32 = —328° F. 

Problems 

1. What centigrade temperature corresponds to the following: 
75° F.? —40° F.? 

2. What Fahrenheit temperature corresponds to the following : 
85° C? 0°CJ —35° C? 

3. Read the Fahrenheit thermometer in your schoolroom, and 
change the reading to the corresponding temperature centigrade. 

4. At what temperature are the readings on both the centi- 
grade and Fahrenheit thermometers the same? 



APPENDIX B 341 

It is evident that the zero of neither the Fahrenheit nor of the 
centigrade scales is the absolute zero of temperature, since tem- 
peratures far below these zeros are often obtained. The calcula- 
tions of scientists have shown that the absolute zero of tempera- 
ture is — 273° C. ; therefore, to change temperature on the centi- 
grade scale to absolute temperatures add 273°. Thus, 27° C. is 
300° absolute temperature, and —200° C. is 73° absolute. 

VII. CHAKLES' LAW 

The law governing the changes of the volume of a gas as its 
temperature changes is known as Charles' law and may be stated 
thus: The volume of a gas varies directly as the absolute tem- 
perature of the gas, the pressure remaining constant. 

The problems connected with this law involve the calculation 
of the new volume assumed by a gas when a change occurs in the 
temperature of the gas. 

Illustrations. — What volume will 600 liters of hydrogen, meas- 
ured at 27° C, have if the temperature drops to 0° C.f 

Solution : 



!7°C. 


= 27° + 273° = 300° absolute 


0°C. 


= 0° + 273° = 273° absolute 



273 

The new temperature (273° abs.) is — of the old temperature; 

273 

hence by Charles' law the new volume is — — of the old volume, 

oUU 

or j^L X S0e-= 273 X 2 = 546 liters. 

U\J%J 

If 1,000 liters of air at —173° C. is heated to 17° C. (the 
pressure remaining constant), what volume will it occupy? 

— 173° C. = 273° — 173° = 100° absolute 
17° C. = 273° + 17° = 290° absolute 

The new temperature (290° absolute) is — of the old tempera- 
ture; hence the new volume is ™9 f the old volume, or, 

io OQO' 

1QQQ X — = 290X10 = 2900 liters. 



342 A PRACTICAL CHEMISTRY 



Problems 

(N. B. — In the following problems the pressure in each case 
is to be taken as standard [760 mm].) 

1. What volume will 970 liters of oxygen measured at 170° 
C. have when the temperature changes to 0° C.f 

2. What volume of oxygen at — 17° C. will fill a 10-liter gas 
holder at 27° C.f 

3. What is the weight of 100 liters of hydrogen measured 
at 33° C.f 

4. Which will weigh the more and how much, 100 liters of 
chlorine measured at 0° C. or 300 liters of air measured at 27° C.f 

5. If a gas holder of 1,000 liters capacity is filled at' 27° C. 
and then cooled to 0° C, what volume of gas will it then contain 1 ? 

6. What volume of acetylene measured at 36° C. will weigh 
5,807 grams? 

7. What weight of sodium chloride will be required to produce 
100 liters of hydrochloric acid measured at 13° C.f 

VIII. BOYLE'S LAW 

Boyle's law may be stated thus: The temperature remaining 
constant, the volume of a gas varies inversely as the pressure to 
which it is subjected. This means that as the pressure increases 
the volume of the gas decreases in the same proportion, and vice 
versa. Calculations involving gas volumes are usually made at 
the standard pressure of 760 mm. of mercury. Gases are often 
measured at pressures other than standard pressure. Hence it 
becomes necessary to calculate the volume which gases measured 
at other pressures would occupy at standard pressure. 

Illustrations. — The temperature remaining constant, 1,000 liters 
of gas measured at 700 mm. pressure will have what volume at 

7 fin 
760 mm. pressure? The new pressure, 760 mm., is — — of the 

old pressure, but the volume varies inversely as the pressure; 

hence the new volume will be -— of the old volume, or 1000 

700 70000 _ 920 + liters. 

760 76 T 



APPENDIX B 343 

Again, What volume of gas measured at 900 mm. will occupy 
1,000 liters volume under standard pressure, the temperature re- 
maining constant? 

The new pressure, 900 mm. is — - of the old pressure, hence the 

*7fifJ 

new volume is — of the old volume, or 
900 

1000 X — = 7600 = 844.4 liters. 



Problems 

(N. B. — In the following problems consider the temperature as 
constant.) 

1. 760 volumes of gas measured at 900 mm. will have what 
volume at 760 mm.? At 1,500 mm.? At 100 mm.? At 1,000 mm.? 

2. A gas holder is filled when the barometer is at 750 mm.; 
if the barometer changes to 770, what fraction of it will be filled ? 
If the barometer changes to 745 mm. what volume of gas must 
escape, provided the capacity of -the holder is 1,000,000 cubic 
feet? 

3. What pressure will be required to compress 1,000,000 cubic 
feet of air, measured under standard conditions, into a tank hav- 
ing a capacity of 250,000 cubic feet? 

It frequently happens that the volume of a gas must be cor- 
rected for both temperature and pressure, i. e., the gas is meas- 
ured at a temperature and pressure other than standard, and its 
volume under standard conditions must be calculated. 

Illustration. — 28 c. c. of gas measured at 17° C. and 750 mm. 
pressure will have what volume under standard conditions? 

17° C. = 17° + 273° = 290° absolute 
0° C. = 0° + 273° = 273° absolute 

273 

The new temperature (273°) is — of the old temperature, 

and the new volume (as corrected for temperature only) will be 

273 

28 X — ! according to the law of Charles. 
290 



344 A PRACTICAL CHEMISTRY 

The new pressure (760 mm.) is — of the old pressure, and 

750 
the new volume (corrected as to pressure only) will be 28 X — ■ 

according to Boyle's law. Correcting for temperature and pres- 

273 750 
sure both, the new volume will be 28 X — X — = 26 01 o c 

290 760 ^°- UiC - c - 

4. 100 c. c. of gas was measured when the thermometer stood 
at 20° C. and the barometer at 755 mm. What volume will the 
gas occupy when the thermometer is at 10° C. and the barometer 
at 765 mm.? 

5. If a rubber balloon be filled with 1,000 liters of hydrogen 
under standard conditions, and then rises to an altitude where 
the barometer stands at 600 mm. and the thermometer at 
— 20° C, what will be the volume of the balloon? What weight 
of zinc will be required to generate enough hydrogen to fill the 
above-mentioned balloon ? 

6. If in the electrolysis of water 100 c. c. of hydrogen meas- 
ured at 20° C. and 765 mm. were generated, what volume of 
oxygen measured under standard conditions was generated at the 
same time in the other tube? What was the weight of this oxy- 
gen? What weight of water was decomposed during the elec- 
trolysis ? 

7. When the barometer is at 760 mm. 50 c. c. of gas is col- 
lected in a tube over a column of mercury 90 mm. high; what 
volume would this gas have at standard pressure? 

IX. AVOGADKO'S HYPOTHESIS 

Under the same conditions of temperature and pressure equal 
volumes of gases contain the same number of molecules. It must 
follow from this that equal numbers of molecules will occupy 
equal volumes. This so-called hypothesis is of value to us in 
several ways: 

(a) In determining the volumes of combining gases when the 
chemical equation is known. 

Illustration. — What volume of oxygen is required to burn 10 
liters of methane? 

CH 4 + 20, = 2H 2 + C0 2 



APPENDIX B 345 

From the equation we see that two molecules of oxygen are 
required to burn one molecule of methane, and since equal num- 
bers of molecules are contained in equal volumes the volume of 
oxygen must be twice as great as the volume of methane, or 20 
liters. 

Problems 

1. What volume of oxygen is required to burn 100 liters of 
hydrosulphuric acid? 

2. How many liters of acetylene can be completely burned 
by using 100 liters of oxygen? 

3. How many liters of products are formed by the complete 
combustion of 100 liters of hydrosulphuric acid? 

(b) In determining the number of atoms in a molecule of an 
elementary gas. 

Illustration. — How many atoms in a molecule of hydrogen? 

By experiment it has been found that one volume of hydrogen 
will combine with one volume of chlorine and form two volumes 
of hydrochloric acid. Now in the two volumes of hydrochloric 
acid there must be (by Avogadro's hypothesis) twice as many 
molecules of hydrochloric acid as there are molecules of hydrogen 
in one volume of hydrogen. Each of these molecules of hydro- 
chloric acid contains an atom of hydrogen. Hence there are 
twice as many atoms of hydrogen as there were molecules of 
hydrogen from which they came; or, in other words, each of the 
molecules of hydrogen must have contained two atoms. 

Problems 

1. It is found by experiment that when two volumes of hy- 
drogen and one volume of oxygen are measured at 100° C. and 
then caused to combine the steam thus formed, if kept at 100° C, 
will occupy two volumes. Calculate how many atoms of oxygen 
are in each molecule of oxygen. 

2. How many atoms of chlorine are in each molecule of 
chlorine ? 

3. If it were possible to substitute pure ozone for oxygen in 
problem 1, it would be found that three volumes of hydrogen and 
one volume of ozone would form three volumes of steam. From 



346 A PRACTICAL CHEMISTRY 

this calculate the number of atoms of oxygen in a molecule of 

ozone. 

(c) In finding the molecular weight of a gas or vapor. 

The molecular weight of a gas or vapor is the weight of the 
molecules of that gas or vapor as compared with some standard, 
either the weight of the hydrogen atom or the one-sixteenth of 
the weight of an oxygen atom. The first step in the process of 
finding the molecular weight by this method is the determination 
of the specific gravity of the gas or vapor as compared with air. 
In the case of gases this is done by the direct weighing of equal 
volumes of the gas and of air; in the case of volatile substances, 
by weighing a definite quantity of the volatile substance, convert- 
ing it into a vapor, measuring the volume of the vapor formed, 
and comparing the weight of the volatile substance with the 
weight of the volume of air equal to the volume of vapor formed. 

Illustrations. — (1) A gas. To find the specific gravity of oxy- 
gen a definite volume of oxygen is weighed and the weight com- 
pared with the weight of an equal volume of air. 

(2) A volatile substance. To find the specific gravity (or 
vapor density) of chloroform vapor. A small quantity of chloro- 
form is carefully weighed and vaporized in such a way that its 
vapor is made to displace air. The displaced air is caught and 
measured, and from the known weight of a liter of air its weight 
is calculated. The weight of the chloroform taken is then com- 
pared with this weight of air (i. e., divided by this weight of air). 
The molecular weight of oxygen is 32, since its molecules contain 
two atoms, and its atom is taken as a standard with a value of 16. 
Air is 0.9046 times as heavy as oxygen. 

When the specific gravity of a gas or the vapor density of a 
compound has been determined the calculation of its molecular 
weight is as follows: 

The product of the vapor density (or specific gravity) of the 
gas or vapor multiplied by 0.9046 is the number of times the 
molecules of the vapor or gas are heavier than oxygen molecules, 
since equal volumes of the vapor and of oxygen contain the same 
number of molecules. This product multiplied by the molecular 
weight of oxygen (32) gives the molecular weight of the gas 
or vapor. 



APPENDIX B 347 

Illustration. — The vapor density of carbon bisulphide is 2.626, 
that is, one volume of its vapor is 2.626 times as heavy as an 
equal volume of air. Air is 0.9046 times as heavy as oxygen. 
2.626 X 0.9046 = 2.375, i. e., carbon bisulphide vapor is 2.375 
times as heavy as oxygen. Its molecules, therefore, are 2.375 
times as heavy as oxygen molecules. Oxygen has a molecular 
weight of 32 ; the molecular weight of carbon bisulphide vapor is 
2.375 X 32 = 76. 

22.4 liters of oxygen under standard conditions weigh 32 grams, 
i. e., as many grams as its molecular weight. Since the weights 
of all gases are proportional to their molecular weights, 22.4 
liters of any gas measured under standard conditions weigh as 
many grams as the molecular weight of the gas. This furnishes a 
means of calculating the molecular weight of a gas when the 
weight of one liter of the gas is known, or, on the other hand, 
of calculating the weight of one liter of the gas when its formula 
is known. 

Problems 

1. Calculate the molecular weight of the compound having 
the vapor density of 2.55. 

2. Alcohol has a vapor density of 1.58; what is its molecular 
weight? 

3. From the molecular weight of wood alcohol (CH 3 OH) cal- 
culate its vapor density. 

4. What is the molecular weight of the compound with a 
vapor density of 0.897? 

X. CALCULATION OF THE FOEMULA OF A COMPOUND 

From the percentage composition of a compound it is an easy 
matter to calculate the simplest formula which will represent the 
compound. 

Illustration. — What is the simplest formula of the compound 
having the composition, C = 92.3 per cent., H = 7.7 per cent. ? 
(1) Divide the percentage of each element by its atomic 
weight : 

C = 92.3 ^ 12 = 7.7 
H . 7.7-t- 1 - 7.7 



348 A PRACTICAL CHEMISTRY 

Therefore C : H : : 7.7 : 7.7 or 1:1, and the simplest formula 
is CH. Whether this is the true formula can be told (if we 
know the molecular weight of the compound) by adding the 
weights of the atoms in this formula and comparing the molec- 
ular weight thus found with the true molecular weight. 

The vapor density of the compound under discussion is 0.897; 
therefore its molecular weight is .897 X -9046 X 32 = 26. 

The molecular weight of CH is 12 -[- 1 = 13. Hence the 
formula CH is not the true formula, and the formula C 2 H 2 , with 
a molecular weight of 26, is the true formula. 

Illustration. — A compound has a vapor density of 3.25, and is 
composed of C = 77.42 per cent., H — 7.52 per cent., N = 15.05 
per cent. Calculate its true formula. 



C = 77.42 
H = 7.52 
N = 15.05 



12 = 6.45 

1 = 7.52 

14 = 1.075 



Reducing to the lowest terms by dividing through by the smallest, 
we have: — C = 6, H = 7, N = 1, or, C 6 H 7 N. The molecular 
weight of the compound calculated from the vapor density = 
3.25 X -9046 X 32 = 93, which corresponds to the molecular 
weight found by adding the weights of the atoms in C 6 H 7 N. 
Hence in this case the simplest formula is also the true formula. 



Problems 
1. Calculate the true formulae of the following: 

(a) C = 92.3%, H = 7 . 7%, vapor density = 2.691 

(b) C = 52.17%, O = 34.78%, H = 13.04%, vapor den- 
sity = 1.58 

(c) O = 76.19%, N = 22.22%, H = 1.58%, vapor 
density = 2.17 

(d) N = 82.35%, H = 17.65%, vapor density = 0.587 

(e) CI = 54.36%, Sn = 45.63%, vapor density = 9 

(f) As = 96.15%, H = 3.84%, vapor density =2.69 



APPENDIX B 349 



XI. FAKADAY'S LAW 



Many facts known to science seem to indicate that the mole- 
cules of acids, bases and salts, when dissolved in water (and 
some other solvents), suffer a process of dissociation or breaking 
to pieces. It has been further found that this dissociation is very- 
slight in concentrated solutions, and increases as the solution is 
diluted. The dissociation is therefore caused by the solvent. It 
has been also found that molecules of the same sort always dis- 
sociate in the same way, i. e., in bases the metal separates from 
the hydroxyl group, while in acids the hydrogen dissociates from 
the other portion. 

When two bodies which are in contact are separated an elec- 
trical charge is developed on each. If the bodies are not in con- 
tact with a conductor of any sort, the charges will remain upon 
them. Water is a non-conductor of electricity; hence when mole- 
cules dissociate their parts remain charged. These charged atoms, 
or groups of atoms, are called ions*. 

Whether ions are positively or negatively charged may be told 
by putting the ends or electrodes of an electric circuit into the 
solution. All positive ions will be attracted to the negative elec- 
trode (or cathode), while the negative ions will be drawn to the 
positive electrode (or anode). Metal and hydrogen ions are thus 
found to be always positively charged. All other ions are nega- 
tively charged. 

During electrolysis the ions are thus attracted to the electrodes. 
When the ions strike the electrodes they lose their electrical 
charges and attack, if possible, the electrodes. Failing in this, 
they may attack the solvent, bubble off as gas, plate the electrodes, 
or be precipitated. Electrolysis, therefore, includes the sorting 
of the ions by the electric current, and also all the various 
changes which they undergo after reaching the electrodes. On 
account of the great variety of results which may be thus obtained, 
electrolysis finds many applications. It is often necessary, there- 
fore, to be able to calculate the quantities of elements liberated 
or deposited or the quantities of the compounds formed. These 
calculations are based on Faraday's law — the quantities of the 



350 A PRACTICAL CHEMISTRY 

various elements liberated by a given electrical current are pro- 
portional to their chemical equivalents. 

By the chemical equivalent of an element is meant the number 
of grams of that element required to combine with or replace one 
gram of a standard element. It is most convenient to use hy- 
drogen as the standard element. The chemical equivalent then 
becomes the number of grams of the element which will combine 
with or replace one gram of hydrogen. Since the valence of 
hydrogen is one, the equivalents of all univalent elements may 
be calculated by dividing their atomic weights by the atomic 
weight of hydrogen. When the valence of the element is more 
than one its equivalent is found by dividing its atomic weight by 
the product of its valence and the atomic weight of hydrogen. 

Illustrations. — The valence of aluminium is 3 ; its atomic weight 

27 1 

is 27.1: therefore its chemical equivalent is equal to ' ftn 

' 3 X 1.008 

or approximately 9. What weight of copper will be deposited 
in an hour by a current which will deposit 100 grams of alu- 
minium in the same length of time? 

Following the plan given above, we find the equivalent of 

copper to be '- z=z 31.5. The chemical equivalent of 

^ F 2 X 1.008 

aluminium (9) : the chemical equivalent of copper (31.5) : : 100 
grams of aluminium : x grams of copper, or, 

9 : 31.5 : : 100 : x 
9x = 3150 
x = 350 grams of copper. 

Problems 

1. During the electrolysis of water how many grams of hy- 
drogen are liberated while one hundred grams of oxygen are 
being liberated ? 

2. In the electrolysis of fused sodium hydroxide how many 
pounds of oxygen will be liberated with 2,305 pounds of sodium? 

3. The cathode in a silver plating bath gains 10 grams of 
silver; how much is lost by the anode? 

4. An electric current can deposit 1,088 grams of silver in a 



APPENDIX B 351 

certain time; how many grams of copper can be deposited by the 
same current in the same length of time? 

5. In the electrolysis of sodium chloride 3,008 grams of 
hydrogen are liberated; how many grams of sodium are liberated 
at the same time? How many grams of chlorine? How many 
grams of sodium hydroxide are formed? 

6. By passing an electric current through a solution of sodium 
sulphate 100 grams of sulphuric acid may be obtained in a certain 
time; what volume of hydrogen measured under standard condi- 
tions will be obtained at the same time? 

7. If a given electric current will liberate 5 grams of hydrogen 
in 20 hours, how long will be required for the same current to 
deposit 55 grams of copper? 

(3) 



INDEX 



Acetates, 160 

Acetic acid, 158, 318 

fermentation, 147, 158 
Acetone, 72 
Acetylene, 76, 307 

burner, 77 

torch, 77 
Acid, 34, 39, 302 

acetic, 158, 318 

boric, 178, 188, 321 

butyric, 149 

capric, 149 

caproic, 149 

carbolic, 176, 186, 187 

carbonic, 63, 168 

chlorplatinic, 263 

citric, 185 

formic, 176 

hydriodic, 184 

hydrobromic, 184 

hydrochloric, 54, 55, 57, 304 

hydrofluoric, 209, 210, 325 

hydrosulphuric, 109, 310 

lactic, 151 

nitric, 113, 131, 311 

nitrous, 119, 125 

oleic, 149 

oxalic, 208, 210 

palmitic, 149 

phosphoric, 99, 145, 309 

picric, 124, 324 



Acid, pyrogallic, 288 

salicylic, 151 

silicic, 214 

stearic, 149 

sulphuric, 105, 110 

sulphurous, 105, 178 

tartaric, 163 

uric, 167 
Acid calcium carbonate, 215 

calcium phosphate, 130, 163 

potassium tartrate, 163 

sodium carbonate, 162 

sodium sulphate, 115 
Agate, 212 
Air, 59, 305 

liquid, 66, 67 
Albumin, 149, 316 
Alcohol, 315 

amyl, 148, 317 

ethyl, 147 

methyl, 176 

test for, 315 

wood, 72, 176 
Alcoholic beverages, 147 

distillation of, 315 

fermentation of, 146, 314 
Aldehydes, 176 
Alkalies, 36 

Alkaloids, 165, 182, 188 
Alloys, 252 
Aluminium, 242, 249 



353 



354 



INDEX 



Aluminium bronze, 253 

chloride, 247 

hydroxide, 246 

oxide, 243 

silicate, 216, 242 

sulphate, 164 
Alums, 163 

Amalgams, 116, 255, 328 
American Cheddar Cheese, 154 
Amethyst, 246 
Ammonia, 116, 125, 312 

fountain, 313 

from coal, 74 

laboratory preparation of, 
117, 312 

water, 116 
Ammonium, 116 

amalgam, 116 

bromide, 288 

chloride, 116 

chlorplatinate, 263 

hydroxide, 116 

molybdate, 309 

nitrate, 120 

nitrite, 305 

phospho-molybdate, 309 

radical, 116 

sulphocyanate, 327 
Amorphous carbon, 71 
Analysis, 1, 3 
Anaesthetics, 179, 188 
Anhydrides, 99 
Animal charcoal, 73 
Anthracene, 198 
Antichlor, 55 
Antidotes, 185, 189 
Anti-friction metals, 257 
Antiseptics, 175, 188 
Antimony, 253 



Antimony sulphide, 100 
Aqua ammonia, 116 

regia, 262 
Argol, 163 
Argon, 65 
Arsenic, 186, 255 

sulphide, 283 
Artificial flavorings, 148 

ice, 118 

perfumes, 148 

silks, 196 
Ash, 72 ' 
Asphaltum, 284 
Atmosphere, 59, 66 
Atomic theory, 41 
Atoms, 41, 47 
Atropine, 183, 186 
Avogadro's hypothesis, 344 

Babbitt's metal, 254 
Babcock's milk tester, 317 
Bacteria, 9, 20, 21, 23, 64, 65, 

134, 151, 176 
Baking powder, 160, 161, 319 

soda, 162 
Banana oil, 148 
Barium chloride, 310 

peroxide, 176 

sulphate, 176 
Barometer, 59, 305 
Bases, 35, 39, 302 
Basic copper carbonate, 252 

lead carbonate, 279 
Bauxite, 243 
Bead, borax, 321 
Beehive oven, 75 
Beer, 148 
Beet sugar, 145 
Bell-metal, 253 



INDEX 



355 



Benzene, 76, 198 
Benzine, 79, 208 
Bessemer converter, 226 

process, 226 
Beverages, 147, 164 
Bichloride of mercury, 177 
Bismuth, 254 
Bituminous coal, 73 
Black gun powder, 122 
Blast furnace, 221 
Bleaching, 54 

powder, 54, 304 
Block tin, 238 

Blowpipe, oxyhydrogen, 26. 32 
Blue prints, 289 
Bone black, 73, 84 
Boric acid, 178, 188, 321 
Borax, 179, 321 
Boyle's Law, 342 
Brandt, 96 
Brandy, 148 
Brass, 252 
Bread, 160, 318 
Bricks, 217, 248 
Britannia metal, 253 
Bromides, 184 
Bromine, 184, 323 
Bronze, 253 

Building materials, 212, 248 
Bunsen burner, 87 
Burettes, 37 
Butter, 149, 156 

fat, 317 
Butter making, 157 
Butyric acid, 149 

Cadmium, 254 
Caffein, 164 
Caleite, 215 



Calcium, 332 

acetate, 158 

borate, 179 

carbide, 77 

carbonate, 63, 136, 215 

chloride, 162 

cyanamide, 135 

fluoride, 209 

hydroxide, 136, 216 

nitrate, 113, 135 

oxide, 136, 216 

phosphate, 97, 145 

silicate, 270 

sulphate, 218 

sulphide, 207 
Calomel, 185, 323 
Candle flame, 87 
Cane sugar, 144 
Canned fruit, 147 
Capric acid, 149 
Caproic acid, 149 
Caramel, 146 
Carats fine, 260 
Carbohydrates, 140, 168 
Carbolic acid, 176, 186, 187 
Carbon, 69 

amorphous, 71 
Carbonate rocks, 215 
Carbonates, 161 
Carbon dioxide, 63, 64, 67, 306 
Carbon disulphide or bisul- 
phide, 110 
Carbonic acid, 63, 168 
Carbon monoxide, 74 
Carbon tetrachloride, 208 
Carborundum, 245, 246 
Carburetter, 83 
Carnallite, 132 
Casein, 154 



356 



INDEX 



Cassiterite, 238 
Castile soap, 204 
Cast iron, 224 
Castner process, 52 
Catalytic agent, 105 
Catalyzer, 105 
Cathartics, 185 
Caustic, lunar, 185 

potash, 185 

soda, 51, 57, 185 
Caustics, 185 
Cellulose, 71, 123, 191 
Cement, 218, 248 
Centigrade thermometer, 6, 296 
Chalk, 215 
Chamber process, 106 
Changes, chemical, 1, 4, 295 

physical, 1, 4, 295 
Charcoal, animal, 73, 84 

filter, 73, 145 

kilns, 72 

wood, 72 
Charles' Law, 341 
Cheese, 154, 171 
Chemical arithmetic, 330 
Chemical change, 1, 4, 47, 295 

properties, 1, 4 

equation, 44, 47 
Chemist, 1 
Chemistry, 1, 4 
Chili saltpeter, 134 
China, 273 

grass, 90, 196 
Chloral, 180 
Chlorides, 54 
Chloride of lime, 54 
Chlorine, 50, 52, 57, 303 
Chloroform, 180, 188, 321 
Chlorplatinic acid, 263 



Chocolate, 165 
Chrome alum, 164 

orange, 283 

red, 283 

yellow, 283 
Chromium, 228 
Cider, 159 

vinegar, 156 
Cinnabar, 177 
Citric acid, 185 
Clabber, 153 • 
Clay, 216, 217 
Cleaning materials, 202 
Coal, 73 

gas, 74, 80 

tar, 74 

tar colors, 198 
Cocaine, 182 
Cocoa, 165 
Codeine, 183 
Coffee, 164 
Coin, gold, 255, 257 

silver, 255, 257 
Coke, 74, 84 
Collodion, 123 
Colmanite, 179 
Combustion, 18, 23, 62, 74 
Compounds, 13, 15, 296 
Concrete, 219 
Contact process, 105 
Converter, Bessemer, 226 
Copper, 240, 249, 267 

alloys, 253 

plating, 266 

sulphate, 242 
Coral, 215 

Corrosive sublimate, 177 
Corn oil, 149 

syrup, 142 



INDEX 



357 



Corundum, 246 
Cotton, 191 

gin, 193 

gun, 123 

seed oil, 149 

wool, 191 
Cream, 156, 171 

separator, 155 
Cream of tartar, 163 
Crucible process, 229 
Crude petroleum, 79 
Cryolite, 243 

Crystallization, 11, 49, 298 
Crystals, 5, 10, 15, 49 
Cupric chloride, 242 

oxide, 242 
Cuprous chloride, 242 

oxide, 242 
Curd, 154 
Cyanide process, 261 

Dalton's atomic theory, 41 
Davy's safety lamp, 93, 94 
Decay, 21, 24 

Decomposition of water, 13, 28 
Definite proportions, 13, 41 
Deliquescence, 12, 15 
Detonators, 124 
Developing, photographic, 288 
Dew, 60 
Dextrin, 144 
Dextrose, 146 
Diamond, 69, 83 
Dioxogen, 176 
Dirt, 202, 205 
Disinfectants, 176 
Distillation, 9, 10 

destructive, 72, 84, 307 
Distilled liquors, 147 



Dust, 65 

Dutch metal, 252 
Dyes, 198, 200, 323 
Dynamite, 124 

Earthenware, 273 
Efflorescence, 12, 15 
Electric furnace, 76, 245 
Electrolysis, 12, 13, 15, 32, 50, 

303 
Electrolytic copper, 241 
Electrotyping, 266 
Elements, 3, 4 

table of, 332 
Emerald, 246 
Emery, 246 
Emetics, 184, 189 
Emetine, 183, 185 
Enamels, 277, 278 
Enzymes, 146 
Epsom salts, 185 
Equations, 44, 47 
Esters, 148 
Etching of glass, 325 
Ethereal salts, 148, 170, 316 
Ethers, 181, 188 
Ethyl acetate, 316 

alcohol, 147 

ether, 180, 321 

sulphuric acid, 181 
Ethylene, 76 
Evaporation, 5 
Explosives, 122 

Fahrenheit thermometer. 6, 296 
Faraday's law, 349 
Fats and oils, 148, 316 
Fehling's solution, 314 
Feldspar, 131, 214 



358 



INDEX 



Fermentation, 146, 169 

acetic, 146, 158 

alcoholic, 146, 314 

lactic, 146 
Ferric ammonium citrate, 289 

chloride, 231 

hydroxide, 185, 327 

nitrate, 230 
Ferriprussiate paper, 289 
Ferrous carbonate, 230 

chloride, 230 

hydroxide, 327 

sulphate, 230 

sulphide, 109 
Fertilizers, 128 
Fibers, 191 
Films, 287 
Filters, charcoal, 73, 145, 306 

water, 10 
Fire damp, 76 

Fire extinguishers, 93, 94, 308 
Flames, 87, 94 

and wire gauze, 92, 308 

Bunsen, 87, 92 

candle, 87 

colored, 89 

luminosity of, 89 

oxidizing, 88, 94 

reducing, 88, 94 

structure of, 87, 307 
Flash powder, 290 
Flavorings, 148 
Flax, 194 
Flint, 212 

glass, 271 
Flowers of sulphur, 102 
Fluorine, 183 
Flux, 224 
Fog, 60 



Food, 140 : 

Formaldehyde, 151, 176, 187, 

320 
Formic acid, 176 
Formulae, 43, 47 
Friction matches, 99 
Fructose, 146 
Fruit sugar, 146 
Fuels, 69 

Fulminate of mercury, 124 
Fuming acids, 169 
Furnaces," blast, 221 

electric, 76, 245 

Hall's aluminium, 242 

puddling, 225 

reverberatory, 233 
Fusel oil, 147 
Fusible metals, 254, 328 

Galena, 233 
Galvanized iron, 240 
Gas, coal, 74, 80 

illuminating, 80, 81 

mantles, 89 

water, 81, 84 
Gases, volume of, 338 
Gasolene, 79, 207 
Gelatin, 149 
German silver, 254 
Glacial acetic acid, 160, 318 
Glass, 268, 277 

etching of, 325 
Glucose, 142, 146, 314 
Glutin, 149 
Glycerine, 124 

trinitrate, 124 
Glyceryl butyrate, 149, 156 

oleate, 149 

palmitate, 149 



INDEX 



359 



Glyceryl stearate, 149 
Gold, 259, 266 

chloride, 262 

mining, 259, 261 

plating, 262 

test for, 328 
Gold schmidt process, 244 
Graphite, 70, 83 
Granite, 131, 214 
Granulated sugar, 144 
Grape sugar, 146 
Gun cotton, 123 

powders, 122 
Gypsum, 218 

Hall's process, 242 
Halogens, 183 
Hardness of water, 205 
Helium, 65 
Hematite, 223 
Hemp, 194 

Hoffman apparatus, 296 
Humidity, 60, 67 
Humus, 136 
Hydraulic main, 80 

mining of gold, 259 
Hydriodic acid, 184 
Hydrobromic acid, 184 
Hydrocarbons, 75, 79, 84 
Hydrochloric acid, 54, 55, 57, 

304 
Hydrofluoric acid, 209, 210, 325 
Hydrogen, 13, 25, 32, 35, 300 

dioxide, 30 

gun, 300 

peroxide, 30, 151, 175, 187 

sulphide, 109 
Hydroquinone, 288 
Hydrosulphuric acid, 109, 310 



Hydroxides, 36 
Hydroxyl, 36 
Hypo, 288 

Iceland spar, 215 
Illuminating gases, 80, 81 
Indicators, 37, 39 
Iodine, 183, 188, 322 
Iodoform, 177, 187, 320 
Ipecac, 185 
Iron, 221, 248 

cast, 224 

compounds, 230 

galvanized, 240 

hydroxides, 231 

oxides, 223 

pyrites, 231 

sulphate, 231 
"sulphide, 231 

wrought, 225, 249 

Japans, 284 
Jamaica ginger, 183 
Jute, 196 

Kainite, 132 

Kaolin, 271 

Kerosene, 8, 79 

Kindling temperature, 18, 23 

Krypton, 65 

Lactic acid, 151 

fermentation, 147, 151 
Lactose, 146 
Lakes, 199, 323 
Lampblack, 283 
Lantern slides, 289 
Lacquer, 279 
Lard, 149 



360 



INDEX 



Laughing gas, 120, 181, 188 

Lavoisier, 17, 23 

Law of definite proportions, 13, 

41 
Law of multiple proportions, 

41 
Lead, 232, 249 

acetate, 237 

carbonate, 279 

chamber process, 106 

chloride, 238 

chromate, 237, 283 

iodide, 238 

nitrate, 237 

oxides, 236 

pencils, 71 

pipe, 235 

sulphate, 237 

sulphide, 233 

white, 237, 279 
Le Blanc process, 206 
Levulose, 146 

Light and chemical action, 287 
Lime, 135, 216 

slacking of, 216, 326 
Limestone, 215 
Lime water, 62 
Limonite, 223 
Linen, 193 
Linseed oil, 273 
Liquid air, 65, 67 
Litharge, 236 
Lithium, 163 
Litmus, 37 
Lunar caustic, 185 



Magnalium, 254 



Magnetite, 223 
Magnesium, 291, 328 



Magnesium citrate, 185 

hydroxide, 185 

sulphate, 185 

sulpho-carbolate, 186 
Manganese dioxide, 20, 53 

sulphate, 53 
Manila hemp, 194 
Mantles, gas, 89 
Manure, 130 
Maple sugar, 145 

syrup, 145 
Marble, 215 
Marsh gas, 75, 307 
Masonry, 217 
Matches, 99 
Medicines, 175 
Mercerizing, 199, 200 
Mercuric chloride, 177, 186, 
187 

oxide, 177 

sulphide, 177 
Mercurous chloride, 185 
Mercury, 177 
Metals, 220 

fusible, 254, 328 
Methane, 75, 76 
Methyl alcohol, 176 

ether, 181 

salicylate, 148 
Metric system, 330 
Milk, 150, 170 

coagulation of, 153, 154 

preserving of, 151, 316 

sugar, 146 
Milk of lime, 216 
Milk of magnesia, 185 
Minerals, 168 
Mixtures, 13, 296 
Molasses, 145 



INDEX 



361 



Molecules, 7 
Molecular weights, 331 
Monopotassium tartrate, 163 
Morphine, 183, 186 
Mortar, 217, 248, 326 
Mother of vinegar, 158 
Mucilage, 148 
Multiple proportions, 41 
Muriatic acid, 56 
Muntz metal, 252 

Naming of acids, 35, 39 

of bases, 36, 39 

of salts, 37, 39 
Naphthalene, 198 
Narcotics, 186 
Nascent state, 43, 47 
Natural gas, 76 
Negatives, photographic, 288 
Neon, 65 

Neutralization, 36, 39, 302 
Nickel-plating, 266 

steel, 227 
Nicotine, 183 
Nitrates, 113 

properties of, 114 

sources of, 113 
Nitric acid, 113, 133, 311 

preparation of, 114, 311 

test for, 312 
Nitric anhydride, 121, 125 

oxide, 115, 120, 313 
Nitrites, 120 
Nitrocellulose, 123 
Nitrogen, 61, 67, 113, 125, 132 

iodide, 314 

pentoxide, 121 

peroxide, 121, 313 

preparation of, 61, 305 



Nitrogen trioxide, 121 
Nitrogenous food, 149. 170 
Nitroglycerine, 123 
Nitrous acid, 119, 125 

anhydride, 121 

oxide, 120, 181, 188 

Oil, cotton seed, 149 

olive, 149 

palm, 149 
Oil of wintergreen, 148 
Oils and fats, 148 
Olein, 149 
Oleic acid, 149 
Open-hearth process, 227 
Opium, 183 
Ores, 221 

Oriental amethyst, 246 
Orpiment, 283 
Osmotic pressure, 326 
Oxalic acid, 208, 210 
Oxidation, 19, 46 
Oxides, 18, 39 

of. nitrogen, 120 

of phosphorus, 99 

of sulphur, 104 
Oxidizing agents, 31, 308 

flames, 88, 94 
Oxygen, 13, 17, 299 

and life, 20, 24 
Oxyhydrogen blowpipe, 26, 

32 
Ozone, 22, 24, 299 

Paints, 279 
Palmitic acid, 149 
Palmitin, 149 
Palm oil, 149 



362 



INDEX 



Paraffin, 79 

oil, 79 

wax, 79 
Paris green, 186 
Parkes' process, 263 
Pasteurizing, 151 
Peat, 73 

Percentage composition, 333 
Perfumes, 148 
Petroleum, 79, 84 
Pewter, 253 
Phenol, 177 
Phenolphthalein, 313 
Phosphates, 96 

in fertilizers, 131 
Phosphine, 99, 309 
Phosphoric acid, 99, 145, 309 
Phosphor bronze, 253 
Phosphorus, 96, 100, 186, 

309 

oxides of, 99 

red, 98 

sesquisulphide, 100 

waxy, 98 

yellow, 98 
Photography, 287 
Physical changes, 1, 4, 295 
Picric acid, 124, 324 
Pig iron, 224 
Pigments, 283, 285 
Pinchbeck, 252 
Plaster, 217, 248 
Plaster of Paris, 217, 248, 

326 
Plate glass, 270 
Plating, 267 

copper, 266 

gold, 262 

nickel, 265 



Plating, silver, 265, 328 
Platinum, 262, 266 

black, 263 
Porcelain, 273 
Potash, 131 

red prussiate of, 231 

yellow prussiate of, 231 
Pot cheese, 154 
Potassium, 129 

acid tartrate, 163 

argenticyani.de, 265 

auricyanide, 262 

bromide, 323 

chlorate, 20 

chloride, 20 

chromate, 283 

chromium alum, 164 

cyanide, 261, 329 

dichromate, 283 

ferricyanide, 231 

ferrocyanide, 231 

hydroxide, 185 

in fertilizers, 132, 311 

iodide, 184 

nitrate, 122 

permanganate, 310 

silicate, 214 

sulphate, 132 
Pottery, 271, 278 
Preserves, 147 
Priestley, Joseph, 17, 23, 

177 
Producer gas, 227 
Proteids, 149 
Protein, 149 
Prussian blue, 232 
Puddling furnace, 225 
Purification of water, 9 
Pyrogallic acid, 288 






INDEX 



363 



Quartz, 212 
Quicklime, 216, 326 
Quinine, 183 

Radicals, 116 
Rain, 6 

Ramie, 90, 196 
Red lead, 236, 284 
Red phosphorus, 98 
Reducing agent, 27, 32 

flame, 88, 94, 308 
Reduction, 46 

by aluminium, 328 

by carbon, 306 

by hydrogen, 28 
Relative humidity, 60, 66 
Rennet, 153, 154 
Respiration, 64 
Retort, 72 
Retort ovens, 75 
Retting, 194 

Reverberatory furnace, 233 
Rock, 212 

crystals, 212 

oil, 79 

salt, 49 
Ruby, 246 

glass, 271 
Rum, 148 

Safety lamp, 93, 94 
Salicylic acid, 151 
Salt, 49, 57 
Salts, 36, 39 
Sand, 212 

Saponification, 203, 210 
Sapphire, 246 
Saturated solutions, 8 
Scrubbers, 81 



Shellac, 284 
Shot, 255 
Siderite, 223 
Silica, 212 
Silicates, 212 
Silicic acid, 214 
Silicon, 212 

bronze, 253 

dioxide, 212 

terrafluoride, 209 
Silk, 196, 200 

artificial, 196 

worm, 196 
Silver, 263, 267 

bromide, 288 

chloride, 287, 289, 304 

coins, -255, 257 

nitrate, 185, 265, 288 

plating, 264, 328 
Sisal, 194 

Slaking of lime, 216, 326 
Slag, 224 
Smithsonite, 239 
Smoke, 18 

Smokeless powder, 123 
Snow, 6, 12 
Soap, 202, 209 

making, 203, 204, 324 
Soda alum, 164 

water, 168 
Sodium, 28, 52 

aluminium alum, 164 

bicarbonate, 318 

carbonate, 206 

chloride, 49, 57 

cj^anide, 52 

hydroxide, 28, 50, 51, 57, 
185 

iodide, 184 



364 



INDEX 



Sodium, nitrate, 113, 134 


Sucrose, 144 


nitrite, 119 


Sugar, 144 


peroxide, 52 


cane, 144 


silicate, 213 


beet, 145 


sulphate, 206 


of lead, 237 


sulphide, 207 


maple, 145 


sulphite, 310 


milk, 146 


thiosulphate, 288 


Suint, 197 


Softening water, 206, 324 


Sulphides, 109, 110, 310 


Soils, 129 


Sulphur, 102, 110, 309 


Solder, 253 


dioxide, 105, 110, 146, 178, 


Solute, 8 


188, 310 


Solution, 7, 8, 14, 297 


flowers, 102 


Solvay process, 162 


occurrence, 102 


Solvent, 8, 324 


plastic, 104 


Specific gravity, 7 


roll, 102 


Speculum metal, 253 


trioxide, 105, 110 


Spiegeleisen, 226 


Sulphuric acid, 105, 110 


Spontaneous combustion, 18, 23 


chamber process, 106 


Spores, 151 


contact process, 105 


Stable manure, 130 


properties of, 108 


Stains, 208, 324 


test for, 310 


Stannic chloride, 327 


Sulphurous acid, 105, 178 


oxide, 239 


Superheater, 83 


Stannous chloride, 327 


Superphosphate of lime, 131 


oxide, 239 


Supersaturated solution, 8 


Starch, 140, 314 


Symbols, 42, 47 


Starch, iodine test for, 142 


Synthesis, 1, 3 


paste, 142 


Syrup, maple, 145 


Steam, 27 




Stearin, 149 


Tallow, 149 


Stearic acid, 149 


Tan bark, 279 


Steel, 225, 249 


Tar, 80 


tempering of, 229 


Tartar, cream of, 163 


Sterilized milk, 152 


Tartaric acid, 163 


Stoneware, 277 


Tea, 164 


Stovaine, 182 


Tellurium, 259 


Stove polish, 71 


Temperature, 6 


Strychnine, 183 


changes, 339 



INDEX 



365 



Tempering steel, 229 
Theobromine, 167 
Theory, atomic, 41 
Thermit, 244 
Thermometers, 6 
Thomas-Gilchrist process, 227 
Tin, 238, 249 

block, 238 

foil, 238 

plate, 238 

roofs, 238 

stone, 238 

ware, 238 
Tinctures, 183, 188 

of arnica, 183 

of iodine, 183, 322 

of iron, 183 

of Jamaica ginger, 183 

of opium, 183 
Toluene, 198 

Toning, photographic, 289 
Tow, 194 

Transparencies, 289 
Tri-iodomethane, 178 
Trinitrophenol, 124 
Turnbull's blue, 232 
Turpentine, 284, 286 
Tuyeres, 221 
Type-metal, 255 

Unsaturated solution, 8 
Uric acid, 167 

Valence, 45, 47 
Varnishes, 284 
Vaseline, 79 
Venetian red, 284 
Ventilation, 64 
Verdigris, 252 



Vermilion, 284 
Vichy, 168, 297 
Vinegar, 156, 171, 318 
mother of, 158, 318 

Washing soda, 207, 210 
Water, 5 

action of, upon sodium, 
28 

composition of, 12, 298 

distilled, 9, 10 

drinking, 8, 15 

filter, 10 

formation of, 27, 29 

hard and soft, 205, 210, 324 

of crystallization, 12, 15, 298 

properties of, 5, 14, 296 

purification of, 9, 297 

softening of, 206, 324 

vapor in the air, 5 

volume changes, 6 
Water-gas process, 82, 83 
Water-glass, 213, 268, 325 
Waxy phosphorus, 98 
Weighting of silk, 197 
Whey, 146, 151 
Whiskey, 148 
White lead, 279 

wash, 216 
Wines, 148 
Wood, 71, 84 

alcohol, 72, 176 

ashes, 72, 132 

charcoal, 72 
Wool, 197 
Wrought iron, 225 

Xenon, 65 
Xylene, 198 



366 
Yeast, 146 



Zinc, 239, 249 
blende, 239 



INDEX 



Zinc chloride, 185 
oxide, 239 
sulphate, 185 

Zincite, 239 

Zymase, 146 



